Understanding the Periodic Table and Trends

6.1 Organizing the Elements Connecting to Your World In 1916, a self-service grocery store opened in Memphis, Tennessee. Shoppers could select items from shelves instead of waiting for a clerk to gather the items for them. In a self- service store, the customers must know how to find the products. From your experience, you know that products are grouped according to sim- ilar characteristics. You don't expect to find fresh fruit with canned fruit, or bottled juice with frozen juice. With a logical clas- sification system, finding and comparing products is easy. In this section, you will learn how elements are arranged in the periodic table and what that arrangement reveals about the elements. Searching For an Organizing Principle A few elements have been known for thousands of years, including copper, silver, and gold. Yet only 13 elements had been identified by the year 1700. Chemists suspected that other elements existed. They had even assigned names to some of these elements, but they were unable to isolate the ele- ments from their compounds. As chemists began to use scientific methods to search for elements, the rate of discovery increased. In one decade (1765-1775), chemists identified five new elements, including three color- less gases-hydrogen, nitrogen, and oxygen. Was there a limit to the num- ber of elements? How would chemists know when they had discovered all the elements? To begin to answer these questions, chemists needed to find a logical way to organize the elements. Chemists used the properties of elements to sort them into groups. In 1829, a German chemist, J. W. Dobereiner (1780-1849), published a clas- sification system. In his system, elements were grouped into triads. A triad is a set of three elements with similar properties. The elements in Figure 6.1 formed one triad. Chlorine, bromine, and iodine may look different. But they have very similar chemical properties. For example, they react easily with metals. Dobereiner noted a pattern in his triads. One element in each triad tended to have properties with values that fell midway between those of the other two elements. For example, the average of the atomic masses of chlorine and iodine is [(35.453 + 126.90)/2] or 81.177 amu. This value is close to the atomic mass of bromine, which is 79.904 amu. Unfortunately, all the known elements could not be grouped into triads. Guide for Reading Key Concepts PARMESANGASTEKKERS • How did chemists begin to organize the known elements? • How did Mendeleev organize his periodic table? • How is the modern periodic table organized? • What are three broad classes of elements? Vocabulary periodic law metals nonmetals metalloid Reading Strategy Comparing and Contrasting As you read, compare and contrast Figures 6.4 and 6.5. How are these two versions of the periodic table similar? How are they different? Figure 6.1 Chlorine, bromine, and iodine have very similar chemical properties. The numbers shown are the average atomic masses for these elements. Chlorine Bromine lodine 35.453 amu 79.904 amu 126.90 amu Section 6.1 Organizing the Elements 155 ГОЛЕТИЕ ПЕРИОДИЧЕСКОГО ЗАКОНА И.МЕНДЕЛЕЕВА Al=27.4 92=6869 Air & 116/13 JOчTA CCCP 1969 OK Figure 6.2 Dimitri Mendeleev proposed a periodic table that could be used to predict the properties of undiscovered elements. Mendeleev's Periodic Table From 1829 to 1869, different systems were proposed, but none of them gained wide acceptance. In 1869, a Russian chemist and teacher, Dmitri Mendeleev, published a table of the elements. Later that year, a German chemist, Lothar Meyer, published a nearly identical table. Mendeleev was given more credit than Meyer because he published his table first and because he was better able to explain its usefulness. The stamp in Figure 6.2 is one of many ways that Mendeleev's work has been honored. Mendeleev developed his table while working on a textbook for his stu- dents. He needed a way to show the relationships among more than 60 ele- ments. He wrote the properties of each element on a separate note card. This approach allowed him to move the cards around until he found an organization that worked. The organization he chose was a periodic table. Elements in a periodic table are arranged into groups based on a set of repeating properties. Mendeleev arranged the elements in his periodic table in order of increasing atomic mass. Figure 6.3 is an early version of Mendeleev's periodic table. Look at the column that starts with Ti = 50. Notice the two question marks between the entries for zinc (Zn) and arsenic (As). Mendeleev left these spaces in his table because he knew that bromine belonged with chlorine and iodine. He predicted that elements would be discovered to fill those spaces, and he predicted what their properties would be based on their locations in the table. The elements between zinc and arsenic were gallium and germa- nium, which were discovered in 1875 and 1886, respectively. There was a close match between the predicted properties and the actual properties of these elements. This match helped convince scientists that Mendeleev's periodic table was a powerful tool. Figure 6.3 In this early version of Mendeleev's periodic table, the rows contain elements with similar properties. Observing A fourth element is grouped with chlorine (CI), bromine (Br), and (I) iodine. What is this element's symbol? 156 Chapter 6 но въ ней, мнѣ кажется, уже ясно выражается примѣнимость вы ставляемаго мною начала ко всей совокупности элементовъ, пай которыхъ извѣстенъ съ достовѣрностію. На этотъ разъ и и желалъ преимущественно найдти общую систему элементовъ. Вотъ этотъ OUNTE Ti-00 V=51 Zr 90 Nb-94 ?=150. Ta-132. Mo=96 W=186. Mu-55 Rh-1044 Pt-197,1 Fe-56 Ru=104, Ir-198. M-C6=59 Pl=106, Os 199. Cu 631 Be=9,1 B-11 Mg-24 Al=27,1 Zu-65,2 Ag=108 Hg=200. Cd=112 2-68 C-12 Si=28 Ur-116 Au=1972 S=118 N=14 P=31 As=75 Sb 122 Bi 210 0-16 S-32 F-19 CI 35,5 Li-7 Na 23 K-39 Se-79,4 Br 80 Rh-85,4 Te=128? 1-127 Cs-133 TI=201 Ca=40 Sr 57,8 Ba=137 Pb=207. 2-45 Ce 92 ?Er-56 La-94 7Y 60 Di=95 ?In 75,6 Th 118? 1 H 2 Li Be 11 12 3 Na Mg 196 20 4 K Ca 37 38 5 Rb Sr Sc Ti 23 24 25 26 27 He 10 B C N O F Ne 13: 14 Al Si P 16 17 18 S Cl Ar 28 29 30 31 32 33 34 35 36 V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 39 40 41 42 43 44 42 43 44 45 46 47 48 49 50 51 52 53 64 Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 72 55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 12 23 74 75 76 77 78 79 n 7 80 81 82 83 84 85 86 6 Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 87 88 89 90 91 92 93 108 109 110 111 112 7 Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Uuu Uub Th Pa U Np Pu 95 96 97 98 99 100 101 102 103 104 105 106 107 1 The Periodic Law The atomic mass of iodine (I) is 126.90. The atomic mass of tellurium (Te) is 127.60. Based on its chemical properties, iodine belongs in a group with bromine and chlorine. So Mendeleev broke his rule and placed tellurium before iodine in his periodic table. He assumed that the atomic masses for iodine and tellurium were incorrect, but they were not. Iodine has a smaller atomic mass than tellurium does. A similar problem occurred with other pairs of elements. The problem wasn't with the atomic masses but with using atomic mass to organize the periodic table. Mendeleev developed his table before scientists knew about the struc- ture of atoms. He didn't know that the atoms of each element contain a unique number of protons. Remember that the number of protons is the atomic number. In 1913, a British physicist, Henry Moseley, determined an atomic number for each known element. Tellurium's atomic number is 52 and iodine's is 53. So it makes sense for iodine to come after tellurium in the periodic table. In the modern periodic table, elements are arranged in order of increasing atomic number. The elements in Figure 6.4 are arranged in order of atomic number, starting with hydrogen, which has atomic number 1. There are seven rows, or periods, in the table. Period 1 has 2 elements, Period 2 has 8 elements, Period 4 has 18 elements, and Period 6 has 32 elements. Each period corre- sponds to a principal energy level. There are more elements in higher num- bered periods because there are more orbitals in higher energy levels. (Recall the rules you studied in Chapter 5 for how electrons fill orbitals.) The elements within a column, or group, in the periodic table have similar properties. The properties of the elements within a period change as you move across a period from left to right. However, the pattern of properties within a period repeats as you move from one period to the next. This pattern gives rise to the periodic law: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties. Checkpoint) How many periods are there in a periodic table? Chec 114 Uug Figure 6.4 In the modern periodic table, the elements are arranged in order of increasing atomic number. Interpreting Diagrams How many elements are there in the second period? Word Origins Periodic comes from the Greek roots peri meaning "around" and hodos, mean- ing "path." In a periodic table, properties repeat from left to right across each period. The Greek word metron means "measure." What does perimeter mean? Section 6.1 Organizing the Elements 157 -≤ -1 1 IA 1A 18 VIIB 8A 13 14 H 2 IIA 2A IIIB IVB VB Metals Metalloids Nonmetals 3A 4A 3 4 Li Be B C 3 4 5 6 7 8 9 10 11 11 12 HIA IVA VA VIA VIA VIIIA 18 22 BUDE Na Mg 38 4B 5B 6B 7B 8B 19 20 21 22 23 24 25 26 27 28 K Ca Sc Ti V Cr Mn Fe Co Ni 37 38 39 40 41 42 43 44 45 46 Rb Sr Y Zr Nb Mo Tc Ru Rh 55 56 71 72 73 74 75 78 77 78 འ༅ལ མ་དང 1B 29 2208 12 13 14 AI Si 30 31 32 Cu Zn Ga Ge 47 48 49 50 Pd Ag Cd In Sn Sb 79 80 81 82 Cs Ba Lu Hf Ta W Re Os Ir Pt 87 88 103 104 105 106 107 108 109 110 111 Au Hg TI 112 Pb C D G 2- 563 15 16 17 IVB VIB 2 SA 6A 7A He 7. 8 10 N F Ne 15 16 17 18 p S CI CLA Ar 33 34 35 36 As Se Br Kr 51 32 54 Te I Xe 83 84 85 86 Bi Po At Rn Fr Ra Lr Rf Db Sg Bh Hs Mt Ds Uuu Uub 114 Uuq 57 58 59 60 61 62 63 64 65 66 67 68 69 70 La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb 89 90 91 92 93 94 95 96 97 98 99 100 101 102 Ac Th Pa U Cf Es Fm Md No Figure 6.5 One way to classify elements in the periodic table is as metals, nonmetals, and metalloids. Inferring What is the purpose for the black stair- step line? Go Online ASTA SC →SCLINKS For: Links on Metals and Nonmetals Visit: www.SciLinks.org Web Code: cdn-1061 Np Pu Am Cm Bk Metals, Nonmetals, and Metalloids Most periodic tables are laid out like the one in Figure 6.5. Some elements from Periods 6 and 7 are placed beneath the table. This arrangement makes the periodic table more compact. It also reflects an underlying structure of the periodic table, which you will study in Section 6.2. Each group in the table in Figure 6.5 has three labels. Scientists in the United States used the labels shown in red. Scientists in Europe used the labels shown in blue. There is some overlap between the systems, but in many cases two differ- ent groups have the same letter and number combination. For scientists to communicate clearly, they need to agree on the stan- dards they will use. The International Union of Pure and Applied Chemis- try (IUPAC) is an organization that sets standards for chemistry. In 1985, IUPAC proposed a new system for labeling groups in the periodic table. They numbered the groups from left to right 1 through 18 (the black labels in Figure 6.5). The large periodic table in Figure 6.9 includes the IUPAC sys- tem and the system used in the United States. The latter system will be most useful when you study how compounds form in Chapters 7 and 8. → Dividing the elements into groups is not the only way to classify them based on their properties. The elements can be grouped into three broad classes based on their general properties. Three classes of elements are metals, nonmetals, and metalloids. Across a period, the properties of ele- ments become less metallic and more nonmetallic. Metals The number of yellow squares in Figure 6.5 shows that most ele- ments are metals-about 80 percent. Metals are good conductors of heat and electric current. A freshly cleaned or cut surface of a metal will have a high luster, or sheen. The sheen is caused by the metal's ability to reflect light. All metals are solids at room temperature, except for mercury (Hg). Many metals are ductile, meaning that they can be drawn into wires. Most metals are malleable, meaning that they can be hammered into thin sheets without breaking. Figure 6.6 shows how the properties of metals can deter- mine how metals are used. 158 Chapter 6 Nonmetals In Figure 6.5, blue is used to identify the nonmetals. These elements are in the upper-right corner of the periodic table. There is a greater variation in physical properties among nonmetals than among metals. Most nonmetals are gases at room temperature, including the main components of air-nitrogen and oxygen. A few are solids, such as sulfur and phosphorus. One nonmetal, bromine, is a dark-red liquid. The variation among nonmetals makes it difficult to describe one set of general properties that will apply to all nonmetals. However, nonmetals are not metals, as their name implies. So they tend to have properties that are opposite to those of metals. In general, nonmetals are poor conductors of heat and electric current. Carbon is an exception to this rule. Solid nonmet- als tend to be brittle, meaning that they will shatter if hit with a hammer. Checkpoint Which type of elements tend to be good conductors of heat and electric current? Iron (Fe) The Gateway Arch in St. Louis, Missouri, is covered in stainless steel containing iron and two other metals, chromium (Cr) and nickel (Ni). The steel is shiny, malleable, and strong. It also resists rusting. Figure 6.6 The metals iron, copper, and aluminum have many important uses. How each metal is used is determined by its properties. Copper (Cu) Copper is ductile and second to only silver as a conductor of electric current. The copper used in electrical cables must be 99.99% pure. 129.** Aluminum (AI) Aluminum is one of the metals that can be shaped into a thin sheet, or foil. To qualify as a foil, a metal must be no thicker than about 0.15 mm. Section 6.1 Organizing the Elements 159 Figure 6.7 Pancake-sized circular slices of silicon, called wafers, are used to make computer chips. Because a tiny speck of dust can ruin a wafer, the people who handle the wafers must wear "bunny" suits. The suits prevent skin, hair, or lint from clothing from entering the room's atmosphere. Metalloids There is a heavy stair-step line in Figure 6.5 that separates the metals from the nonmetals. Most of the elements that border this line are shaded green. These elements are metalloids. A metalloid generally has properties that are similar to those of metals and nonmetals. Under some conditions, a metalloid may behave like a metal. Under other conditions, it may behave like a nonmetal. The behavior often can be controlled by changing the conditions. For example, pure silicon is a poor conductor of electric current, like most nonmetals. But if a small amount of boron is mixed with silicon, the mixture is a good conductor of electric current, like most metals. Silicon can be cut into wafers, like those being inspected in Figure 6.7, and used to make computer chips. 1. 2. 3. 4. 6.1 Section Assessment Key Concept How did chemists begin the process of organizing elements? Key Concept What property did Mendeleev use to organize his periodic table? Key Concept How are elements arranged in the modern periodic table? Key Concept Name the three broad classes of elements. 5. Which of these sets of elements have similar physical and chemical properties? a. oxygen, nitrogen, carbon, boron b. strontium, magnesium, calcium, beryllium c. nitrogen, neon, nickel, niobium 6. Identify each element as a metal, metalloid, or nonmetal. a. gold c. sulfur 160 Chapter 6 b. silicon d. barium 7. Name two elements that have properties similar to those of the element sodium. Connecting Concepts Atomic Number What does an atomic number tell you about the atoms of an element? Why is atomic number better than atomic mass for organizing the elements in a periodic table? Use what you learned in Section 4.2 to answer this question. Inter teractive Textbook Assessment 6.1 Test yourself on the concepts in Section 6.1. with ChemASAP 6.2 Classifying the Elements Connecting to Your World The sculptor Augustus Saint- Gaudens designed this gold coin at the request of Theodore Roosevelt. President Roosevelt wanted coins minted in the United States to be as beautiful as ancient Greek coins, which he admired.The coin is an example of a double eagle. The name derives from the fact that the coin was worth twice as much as $10 coins called eagles. A coin may contain a lot of information in a small space-its value, the year it was minted, and its country of origin. Each square in a peri- odic table also contains a lot of informa- tion. In this section, you will learn what types of information are usually listed in a periodic table. Squares in the Periodic Table The periodic table displays the symbols and names of the elements, along with information about the structure of their atoms. Figure 6.8 shows one square from the detailed periodic table of the elements in Figure 6.9 on page 162. In the center of the square is the symbol for sodium (Na). The atomic number for sodium (11) is above the symbol. The element name and average atomic mass are below the symbol. There is also a verti- cal column with the numbers 2, 8, and 1, which are the number of electrons in each occupied energy level of a sodium atom. The symbol for sodium is printed in black because sodium is a solid at room temperature. In Figure 6.9, the symbols for gases are in red. The sym- bols for the two elements that are liquids at room temperature, mercury and bromine, are in blue. The symbols for some elements in Figure 6.9 are printed in green. These elements are not found in nature. In Chapter 25, you will learn how scientists produce these elements. The background colors in the squares are used to distinguish groups of elements. For example, two shades of gold are used for the metals in Groups 1A and 2A. The Group 1A elements are called alkali metals, and the Group 2A elements are called alkaline earth metals. The name alkali comes from the Arabic al aqali, meaning "the ashes." Wood ashes are rich in com- pounds of the alkali metals sodium and potassium. Some groups of non- metals also have special names. The nonmetals of Group 7A are called halogens. The name halogen comes from the combination of the Greek word hals, meaning salt, and the Latin word genesis, meaning "to be born." There is a general class of compounds called salts, which include the com- pound called table salt. Chlorine, bromine and iodine, the most common halogens, can be prepared from their salts. Guide for Reading Key Concepts • What type of information can be displayed in a periodic table? • How can elements be classified based on their electron configurations? Vocabulary alkali metals alkaline earth metals halogens noble gases representative elements transition metal inner transition metal Reading Strategy Relating Text and Visuals As you read, look carefully at Figure 6,9. After you read the section, explain what you can tell about an element from the color assigned to its square and the color assigned to its symbol. Figure 6.8 This is the element square for sodium from the periodic table in Figure 6.9. Interpreting Diagrams What does the data in the square tell about the structure of sodium atoms? you 11 Na Sodium 22.990, -Atomic number Electrons in each energy level Element symbol -Element name Average atomic mass Section 6.2 Classifying the Elements 161 1 1A 1 H Hydrogen 1.0079 1 2 2A 2 2 N N Periodic Table of the Elements Representative Elements Transition Elements Alkali Metals Transition Metals C Solid Alkaline Earth Metals Inner transition metals Other Metals Br Liquid Hel Gas Metalloids Li Lithium 6.941 11 Be Beryllium 9.0122 12 Na Mg Sodium 22.990 19 K Potassium 39.098 37 Magnesium 24.305 20 Ca Calcium 40.08 Nonmetals Noble Gases Not found Te in nature 2 NJ 00 NJ 8 2 3 4 5 6 7 N 00 00 N 8 2 24 0 0 24 2 3B 21 Sc Scandium 22 Ti Titanium 44.9564790 8 10 4B 5B 6B 7B 9 8B 8 $ 2002 23 V na 23.00 ho 24 25 26 27 Vanadium 50.941 2 2 2 38 8 39 8 40 41 8. 42 18 19 18 18 18 18 Rb 8 1 Sr મૈં Y 9 Zr 10 2 Nb 12 1 Mo 13 17: Cr Chromium 51.996 Manganese 54.938 13 15 Mn2 Fe Co NGON 2 Iron 55.847 Cobalt 58.933 2883. 2 43 18. 14. - $0.00 44 45 Ru 18 15 18 Rh 16 1 Rubidium 85.468 55 Cs Cesium 132.91 18 Strontium 87.62 Yttrium 88.906 2 18 100 00 00 00 - 56 8 71 Zirconium 91.22 272 Niobium 92.906 સ 8 Molybdenum 95.94 74 2 Technetium (98) Ruthenium 101:07: Rhodium 102.91: 75 ? 76 77 2 18 18 18 18 8 Ba 18 મ Lu 32 Hf 10 32 Ta 32 17 W 18 18 18 18 32 12 Re 32 13 Os 32 14 Ir 32 15 2 2 2. 2 2 2 NONGON 8 2 Barium 137.33 Lutetium 174.97 Hafnium 178.49 Tantalum 180.95 Tungsten 183.85 Rhenium 186.21. Osmium 190.2 Iridium 192.22 2 87 8 88 8 103 104 105 106 107 108 18 18 18 18 18 18: Fr 32 18 Ra 32 32. 32 18 32 32 Db 32 32 32 32 Sg Bh Hs 32 32 8 8 9 10 Francium 1 (223) Radium (226) 2 2 Lawrencium Rutherfordium ን (262 (261) Dubnium (262) 11 2 32 13 32 Seaborgium 2 (263) Bohrium 2 (264) 14 Hassium 2 (265) NANNOIN 8 109 18 Mt 32 32 15 Meitnerium 2 200224 N 8 (268) Lanthanide Series 57 La 58 59 60 61 62 19 18 18 Ce 20 Pr Nd Sm 24 Lanthanum 138.91 Cerium 140.12 Praseodymium 140.91 Neodymium 144.24 Promethium (145) Samarium 150.4 Actinide Series 89 90 91 92 93 94 Ac Th 19 32 Pa 321 20 U 32 32 24 10 Actinium N Thorium 232.04 Protactinium 231 04 Uranium 238.03 Neptunium (237) Plutonium (244) Figure 6.9 In this periodic table, the colors of the boxes are used to classify representative elements and transition elements. (227) 162 Chapter 6 Atomic number 14 Electrons in each energy level Si Element symbol 13 14 15 16 17 Silicon -Element name 3A 44 5A 6A 7A 18 8A 2 He Hehur 4.0028 28.086 5 6 7 8 9 10 Average atomic mass B C N 0 F Ne Boron 10.81 13 Carbon 12.011 Nitrogen 14.007 Oxygen 15.999 Fluorine 18.998 Neon 20.179 200 8 14 3 ΑΙ Si P UG 15 evo in 16 6. 10 11 12 NOON. 28 Ni Nickel 58.71 46 36 18: 18 Pd 18 Palladium 106.4 78 Pt Platinum 195.09 1B 29 Cu Copper 63.546 47 18 Zinc 65.38 48 Ag Cd Silver 107.87 79 Cadmium 112.41 18 18 18 Gallium 69.72 49 In Indium 114.82 81 TI 30 18 Zn ཀྑ ན བ ༦ 2B Aluminum 26.982 Silicon 28.088 Phosphorus 30.974 S Sulfur 32.06 17 CI Chlorine 35.453 Argon Ar 2882 2 31 4 32 33 18 Ga 3 As 04 00 00 00 102 2 2 8 51 52 53 18 18 18 18 3 Sn 18 Te I 18 Tin 118.89 Ge Germanium 72-59 50 *4 00 00 00 *4* Arsenic 74.922 Sb Antimony 12175 Selenium 78.96 35 Br Bromine 79.904 18 28888 36 Kr Krypton lodine 126.90 Xe Xenon 13130 34 18 Se 70000 60 2007 Tellurium 127.60 80 18 18 32 Au 32 18 Hg 28832183 2 2 2 82 B 83 8 84 8 85 36 18 18 18 Pb 32 18 Bi 32 18 Po 32 18 At Rn 4 $ 6 Gold 196.97 Mercury 200.59 Thallium Lead Bismuth 204.37 207.2 208.98 Polonium (209) Astatine (210) Radon 110 8 18 32 32 NOONGE 111 32 32 18 CONNO CON 112 114 Uub 18 32 32 18 Darmstadtium 1 (269) Unununium (272) Ununbium (277) *Uuq Ununquadium *Name not officially assigned. 63 Eu Europium 151.96 Gadolinium 157.25 64 65 66 67 63 69 70 Gd 25 Tb Dy Ho Er Tm Yb Terbium 358.93 Dysprosium 162.50 Holmium 164.93 Erbium 167.26 Thulium 168.93 Ytterbium 173.04 95 96 97 98 99 100 101 102 30 Americium (243) Curium (247) Berkelium (247) Californium (251) Einsteinium (252) Fermium (257) Mendelevium (258) Nobelium 12591 Section 6.2 Classifying the Elements 163 Figure 6.10 This blimp contains helium, one of the noble gases. Applying Concepts What does the ability of a helium- filled blimp to rise in air tell you about the density of helium? Go Online AASTE SCLINKS For: Links on Chemical Families Visit: www.SciLinks.org Web Code: cdn-1062 Electron Configurations in Groups Electrons play a key role in determining the properties of elements. So there should be a connection between an element's electron configuration and its location in the periodic table. Elements can be sorted into noble gases, representative elements, transition metals, or inner transition metals based on their electron configurations. You may want to refer to Figure 6.9 as you read about these classes of elements. The Noble Gases The blimp in Figure 6.10 is filled with helium. Helium is an example of a noble gas. The noble gases are the clements in Group 9A of the periodic table. These nonmetals are sometimes called the inert gases because they rarely take part in a reaction. The electron configurations for the first four noble gases in Group 8A are listed below. Helium (He) Neon (Ne) Argon (Ar) Krypton (Kr) 152 1s22s22p 1s22s22p 3s23p 1s22s22p63s23p63d1o4s24p6 Look at the description of the highest occupied energy level for each ele- ment, which is highlighted in yellow. The s and p sublevels are completely filled with electrons. Chapter 7 will explain how this arrangement of elec- trons is related to the relative inactivity of the noble gases. The Representative Elements Figure 6.11 shows the portion of the periodic table containing Groups 1A through 7A. Elements in these groups are often referred to as representative elements because they display a wide range of physical and chemical properties. Some are metals, some are non- metals, and some are metalloids. Most of them are solids, but a few are gases at room temperature, and one, bromine, is a liquid. In atoms of representative elements, the s and p sublevels of the high- est occupied energy level are not filled. Look at the electron configurations for lithium, sodium, and potassium. In atoms of these Group 1A elements, there is only one electron in the highest occupied energy level. The electron is in an s sublevel. Lithium (Li) Sodium (Na) Potassium (K) 152251 1522522p63s1 1s22s22p3s23p°45' In atoms of carbon, silicon, and germanium, in Group 4A, there are four electrons in the highest occupied energy level. Carbon (C) Silicon (Si) Germanium (Ge) 1s22s22p2 1s22522p63523p2 1s22s22p63523p 3d1o4s24p2 For any representative element, its group number equals the number of electrons in the highest occupied energy level. Check Checkpoint) Why are noble gases sometimes referred to as inert gases? 164 Chapter 6 Magnesium This magnified view of a leaf shows the green structures where light energy is changed into chemical energy. The compound chlorophyll, which contains magnesium, absorbs the light. Sodium When salt lakes evaporate, they form salt pans like this one in Death. Valley, California. The main salt in a salt pan is sodium chloride. Figure 6.11 Some of the representative elements exist in nature as elements. Others are found only ' in compounds. Arsenic This bright red. ore is a major source of arsenic in Earth's crust. It contains a compound of arsenic and sulfur. H Hydrogen BA 8 Li Be B C N 9.0122 - 12 Lithium 4941 11 Beryllium Na Mg Godium 72960 Magnesium 24.305 Aluminum 26.062 Porad 1051 Carbon 12.011 Nitrogen 14,002 Oxygen 15.599 Florine 18.598 13 15 16. 17. AI Si P S CI Silicon Phosphorus Sillfor Chlorine: 30.974 32.06 38.953 * 13 .20 31 34 35 B Patessturn K Ca Calcium 40.08 Gallium 60.72 Ga Ge As Se Hin 74922 Selenium 78.06 Bronine 79.204 2 38 49 50 » 53 Rb Sr In Sn Sb Te I Sariu 137.33 Rukidhan Strontium 87.62 indium 114.82 Tin 118.69 55 56 B1 話 82 Cs Ba TI Pb Cesium Thallium Lead 13201 204,37 207.2 88 Fr Ra Fransdam (223) Azdium (220) Anthony 2175 52200 lodine 126.90 * 1 * * 83 84 85 Bi 73 Po At * Bismuth 208.92 Poloniuni (209) Astatine 2210 Sulfur These scientists are sampling gases being released from a volcano through a vent called a fumarole. The yellow substance is sulfur. Transition Elements In the periodic table, the B groups separate the A groups on the left side of the table from the A groups on the right side. Elements in the B groups, which provide a connection between the two sets of representative ele- ments, are referred to as transition elements. There are two types of transi- tion elements-transition metals and inner transition metals. They are classified based on their electron configurations. The transition metals are the Group B elements that are usually dis- played in the main body of a periodic table. Copper, silver, gold, and iron are transition metals. In atoms of a transition metal, the highest occupied s sublevel and a nearby d sublevel contain electrons. These elements are characterized by the presence of electrons in d orbitals. The inner transition metals appear below the main body of the peri- odic table. In atoms of an inner transition metal, the highest occupied s sublevel and a nearby ƒ sublevel generally contain electrons. The inner transition metals are characterized by ƒ orbitals that contain electrons. Before scientists knew much about inner transition metals, people began to refer to them as rare-earth elements. This name is misleading because some inner transition metals are more abundant than other elements. Blocks of Elements If you consider both the electron configurations and the positions of the elements in the periodic table, another pattern emerges. In Figure 6.12, the periodic table is divided into sections, or blocks, that correspond to the highest occupied sublevels. The s block con- tains the elements in Groups 1A and 2A and the noble gas helium. The p block contains the elements in Groups 3A, 4A, 5A, 6A, 7A, and 8A, with the exception of helium. The transition metals belong to the d block, and the inner transition metals belong to the ƒ block. You can use Figure 6.12 to help determine electron configurations of elements. Each period on the periodic table corresponds to a principal energy level. Say an element is located in period 3. You know that the sand p sublevels in energy levels 1 and 2 are filled with electrons. You read across period 3 from left to right to complete the configuration. For transition ele- ments, electrons are added to a d sublevel with a principal energy level that is one less than the period number. For the inner transition metals, the principal energy level of the fsublevel is two less than the period number. This procedure gives the correct electron configurations for most atoms. Figure 6.12 This diagram classifies elements into blocks according to sublevels that are filled or filling with electrons. Interpreting Diagrams In the highest occupied energy level of a halogen atom, how many electrons are in the p sublevel? s block p block d1 d2 d3 d4 d5 do do do do dio d block fblock f2 f3 f4 f5 fô f7 få fŷ f10 f11 f12f13f14 166 Chapter 6 CONCEPTUAL PROBLEM 6.1 Using Energy Sublevels to Write Electron Configurations Nitrogen is an element that organisms need to remain healthy. How- ever, most organisms cannot obtain nitrogen directly from air. A few organisms can convert elemental nitrogen into a form that can be absorbed by plant and animal cells. These include bacteria that live in lumps called nodules on the roots of legumes. The photograph shows the nodules on a bean plant. Use Figure 6.12 to write the elec- tron configuration for nitrogen (N), which has atomic number 7. Analyze Identify the relevant concepts. For all elements, the atomic number is equal to the total number of electrons. For a representa- tive element, the highest occupied energy level is the same as the number of the period in which the element is located. From the group in which the element is located, you can tell how many electrons are in this energy level. Practice Problems b. strontium 8. Use Figure 6.9 and Figure 6.12 to write the elec- tron configurations of the following elements. a. carbon c. vanadium (Hint: Remember that the principal energy level number for elements in the d block is always one less than the period number.) Solve Apply concepts to this situation. Nitrogen is located in the second period of the periodic table and in the third group of the p block. Nitrogen has seven electrons. Based on Figure 6.12, the configuration for the two elec- trons in the first energy level is 1s2. The config- uration for the five electrons in the second energy level is 2s 22p3. 9. List the symbols for all the elements whose electron configurations end as follows. Each n represents an energy level. a. ns2np1 b. ns2nps c. ns2npond2(n+1)s2 6.2 Section Assessment 10. Key Concept What information can be included in a periodic table? 11. Key Concept Into what four classes can elements be sorted based on their electron configurations? 12. Why do the elements potassium and sodium have similar chemical properties? 13. Classify each element as a representative element, transition metal, or noble gas. a. 1s22s22p°3s23p63d1o4s24p b. 1s22s22p63s23p®3d®4s2 c. 1s22s22po3s23p2 14. Which of the following elements are transition metals: Cu, Sr, Cd, Au, Al, Ge, Co? 15. How many electrons are in the highest occupied energy level of a Group 5A element? Elements Handbook Noble Gases Look at the atomic properties of noble gases on page R36. Use what you know about the structure of atoms to explain why the color produced in a gas discharge tube is different for each gas. Interactive Textbook Assessment 6.2 Test yourself on the concepts in Section 6.2. with ChemASAP Section 6.2 Classifying the Elements 167 Technology & Society True Colors Paint consists essentially of a pigment, a binder, and a liquid in which the other components are dissolved or dispersed. The liquid keeps the mixture thin enough to flow. The binder attaches the paint to the surface being painted, and the pigment determines the color. Pigments may be natural or manufactured. They may be inorganic or organic. The same pigment can be used in a water-based or oil-based paint. Comparing and Contrasting Describe at least three differences between the cave painting and the painting by Jacob Lawrence. Yellow ochre 168 Chapter 6 Natural pigments A prehistoric artist had a limited choice of colors- black from charcoal and red, brown, and yellow from oxides of iron in Earth's crust.These oxides (or ochre) pigments are often referred to as earth tones. Prehistoric art Around 14,000 years ago, an artist painted this bison on the ceiling of a cave in Spain. It is about two meters long. Charred wood is a source of charcoal, Red ochre Manganese Violet Red Iron Oxide Cadmium Orange M Cobalt Blue Cobalt Yellow Zinc White Chromium Oxide Green Manufactured pigments Alchemists (and then chemists) made pigments that don't exist in nature. They also made purer versions of natural pigments. Many of these pigments contain transition metals. From pigments to paint Artists mixed manufactured pigments with binders and solvents to make paint. Although premixed paints became available around 1800, some artists, including Jacob Lawrence, continued to mix their own paints. The Builders, 1974, by Jacob Lawrence M 6.3 Periodic Trends Guide for Reading Key Concepts • What are the trends among the elements for atomic size? • How do ions form? • What are the trends among the elements for first ionization energy, ionic size, and electronegativity? • What is the underlying cause of periodic trends? Vocabulary atomic radius ion cation Connecting to Your World An atom doesn't have a sharply defined boundary. So the radius of an atom cannot be measured directly. There are ways to estimate the sizes of atoms. In one method, a solid is bombarded with X rays, and the paths of the X rays are recorded on film. Sodium chloride (table salt) produced the geometric pattern in the photograph.Such a pattern can be used to calculate the position of nuclei in a solid. The distances between nuclei in a solid are an indication of the size of the parti- cles in the solid. In this section, you will learn how properties such as atomic size are related to the location of elements in the periodic table. anion ionization energy electronegativity Reading Strategy Building Vocabulary After you read this section, explain the difference between a cation and an anion. Figure 6.13 This diagram lists the atomic radii of seven nonmetals. An atomic radius is half the distance between the nuclei of two atoms of the same element when the atoms are joined. Trends in Atomic Size Another way to think about atomic size is to look at the units that form when atoms of the same element are joined to one another. These units are called molecules. Figure 6.13 shows models of molecules (molecular models) for seven nonmetals. Because the atoms in each molecule are identical, the distance between the nuclei of these atoms can be used to estimate the size of the atoms. This size is expressed as an atomic radius. The atomic radius is one half of the distance between the nuclei of two atoms of the same element when the atoms are joined. The distances between atoms in a molecule are extremely small. So the atomic radius is often measured in picometers. Recall that there are one trillion, or 1012, picometers in a meter. The molecular model of iodine in Figure 6.13 is the largest. The distance between the nuclei in an iodine mol- ecule is 280 pm. Because the atomic radius is one half the distance between the nuclei, a value of 140 pm (280/2) is assigned as the radius of the iodine atom. In general, atomic size increases from top to bottom within a group and decreases from left to right across a period. Distance between nuclei Nucleus Hydrogen (H2) 30 pm Oxygen (O2) Nitrogen (N2) 66 pm 70 pm Atomic radius Fluorine (F2) 62 pm Chlorine (Cl2) 102 pm Bromine (Br2) 120 pm lodine (12) 140 pm 170 Chapter 6 Atomic radius (pm) Atomic Radius Versus Atomic Number 300 Period 4 Period 5 Cs Period 3 Rb Period 2 250 200 150 Period Li Na 100 g 50 He Sc Cd Zn 0 10 20 30 40 50 60 Atomic number Group Trends in Atomic Size In the Figure 6.14 graph, atomic radius is plotted versus atomic number. Look at the data for the alkali metals. and noble gases. The atomic radius within these groups increases as the atomic number increases. This increase is an example of a trend. As the atomic number increases within a group, the charge on the nucleus increases and the number of occupied energy levels increases. These variables affect atomic size in opposite ways. The increase in positive charge draws electrons closer to the nucleus. The increase in the number of occupied orbitals shields electrons in the highest occupied energy level from the attraction of protons in the nucleus. The shielding effect is greater than the effect of the increase in nuclear charge. So the atomic size increases. Periodic Trends in Atomic Size Look again at Figure 6.14. In general, atomic size decreases across a period from left to right. Each element has one more proton and one more electron than the preceding element. Across a period, the electrons are added to the same principal energy level. The shielding effect is constant for all the elements in a period. The increas- ing nuclear charge pulls the electrons in the highest occupied energy level closer to the nucleus and the atomic size decreases. Figure 6.15 summa- rizes the group and period trends in atomic size. Trends in Atomic Size Size generally decreases Size generally increases Figure 6.14 This graph plots atomic radius versus atomic number for 55 elements. INTERPRETING GRAPHS a. Analyzing Data Which alkali metal has an atomic radius of 238 pm? b. Drawing Conclusions Based on the data for alkali metals and noble gases, how does atomic size change within a group? c. Predicting is an atom of barium, atomic number 56, smaller or larger than an atom of cesium (Cs)? Figure 6.15 The size of atoms tends to decrease from left to right across a period and increase from top to bottom within a group. Predicting If a halogen and an alkali metal are in the same period, which one will have the larger radius? Section 6.3 Periodic Trends 171 Figure 6.16 When a sodium Nucleus 11 p* 12 n° 11 e Sodium atom (Na) atom loses an electron, it Nucleus becomes a positively charged 17 p* ion. When a chlorine atom 18 по 17 gains an electron, it becomes a negatively charged ion. Interpreting Diagrams What happens to the protons and neutrons during this change? lons Chlorine atom (CI) Lose one electron -1e- ر Gain one electron +1e 10 e Nucleus + 11 pt 12 n° Sodium ion (Na') 'Nucleus 17 p* 18 @ 18h0 Chloride ion (CITM) teractive Textbook Animation 7 Discover the ways that atoms of elements combine to form compounds. with ChemASAP Some compounds are composed of particles called ions. An ion is an atom or group of atoms that has a positive or negative charge. An atom is electri- cally neutral because it has equal numbers of protons and electrons. For example, an atom of sodium (Na) has 11 positively charged protons and 11 negatively charged electrons. The net charge on a sodium atom is zero [(11+) + (−−11) = 0]. Positive and negative ions form when electrons are transferred between atoms. Atoms of metallic elements, such as sodium, tend to form ions by losing one or more electrons from their highest occupied energy levels. A sodium atom tends to lose one electron. Figure 6.16 compares the atomic structure of a sodium atom and a sodium ion. In the sodium ion, the number of electrons (10) is no longer equal to the number of protons (11). Because there are more positively charged protons than negatively charged electrons, the sodium ion has a net positive charge. An ion with a positive charge is called a cation. The charge for a cation is written as a number followed by a plus sign. If the charge is 1+, the number 1 is usually omitted from the complete symbol for the ion. So Na* is equivalent to Na1*. Atoms of nonmetallic elements, such as chlorine, tend to form ions by gaining one or more electrons. A chlorine atom tends to gain one electron. Figure 6.16 compares the atomic structure of a chlorine atom and a chlo- ride ion. In a chloride ion, the number of electrons (18) is no longer equal to the number of protons (17). Because there are more negatively charged electrons than positively charged protons, the chloride ion has a net nega- tive charge. An ion with a negative charge is called an anion. The charge for an anion is written as a number followed by a minus sign. Check Checkpoint) What is the difference between a cation and an anion? 172 Chapter 6 Trends in lonization Energy Recall that electrons can move to higher energy levels when atoms absorb energy. Sometimes there is enough energy to overcome the attraction of the protons in the nucleus. The energy required to remove an electron from an atom is called ionization energy. This energy is measured when an element is in its gaseous state. The energy required to remove the first electron from an atom is called the first ionization energy. The cation produced has a 1+ charge. First ionization energy tends to decrease from top to bottom within a group and increase from left to right across a period. Table 6.1 lists the first, second, and third ionization energies for the first 20 elements. The second ionization energy is the energy required to remove an electron from an ion with a 1+ charge. The ion produced has a 2+ charge. The third ionization energy is the energy required to remove an electron from an ion with a 2+ charge. The ion produced has a 3+ charge. Ionization energy can help you predict what ions elements will form. Look at the data in Table 6.1 for lithium (Li), sodium (Na), and potassium (K). The increase in energy between the first and second ionization ener- gies is large. It is relatively easy to remove one electron from a Group 1A metal atom, but it is difficult to remove a second electron. So Group 1A metals tend to form ions with a 1+ charge. Table 6.1 Ionization Energies of First 20 Elements (kJ/mol) Symbol First Second Third H 1312 He (noble gas) 2372 5247 Li 520 7297 11,810 Be 899 1757 14,840 B C 801 2430 3659 1086 2352 4619 N 1402 2857 4577 1314 3391 5301 F 1681 3375 6045 Ne (noble gas) 2080 3963 6276 Na 496 4565 6912 Mg 738 1450 7732 ΑΙ 578 1816 2744 Si 786 1577 3229 P 1012 1896 2910 S 999 2260 3380 K 8 지곡으 1256 2297 3850 Ar (noble gas) 1520 2665 3947 419 3069 4600 Ca 590 1146 4941 *An amount of matter equal to the atomic mass in grams. Section 6.3 Periodic Trends 173 Figure 6.17 This graph reveals. group and period trends for ionization energy. INTERPRETING GRAPHS a. Analyzing Data Which element in period 2 has the lowest first ionization energy? In period 3? b. Drawing Conclusions What is the group trend for first ionization energy for noble gases and alkali metals? c. Predicting If you drew a graph for second ionization energy, which element would you have to omit? Explain. terary.wwwAN, AARON She is gun wwww www.tw. EAT ADALA Figure 6.18 First ionization energy tends to increase from left to right across a period and decrease from top to bottom within a group. Predicting Which element would have the larger first ionization energy-an alkali metal in period 2 or an alkali metal in period 4? 174 Chapter 6 First ionization energy (kJ/mol) First lonization Energy Versus Atomic Number 2500 He Ne 2000 1500 1000 N Xe Be Zn As Cd Mg 500 Li Na K Rb CS 0 10 20 30 40 50 60 Atomic number Group Trends in lonization Energy Figure 6.17 is a graph of first ion- ization energy versus atomic number. Each red dot represents the data for one element. Look at the data for the noble gases and the alkali metals. In general, first ionization energy decreases from top to bottom within a group. Recall that the atomic size increases as the atomic number increases within a group. As the size of the atom increases, nuclear charge has a smaller effect on the electrons in the highest occupied energy level. So less energy is required to remove an electron from this energy level and the first ionization energy is lower. Periodic Trends in lonization Energy In general, the first ionization energy of representative elements tends to increase from left to right across a period. This trend can be explained by the nuclear charge, which increases, and the shielding effect, which remains constant. The nuclear charge increases across the period, but the shielding effect remains con- stant. So there is an increase in the attraction of the nucleus for an electron. Thus, it takes more energy to remove an electron from an atom. Figure 6.18 summarizes the group and period trends for first ionization energy. Trends in First lonization Energy Energy generally increases Energy generally decreases Quick LAB Periodic Trends in lonic Radii Purpose Make a graph of ionic radius versus atomic number and use the graph to identify periodic and group trends. Materials • graph paper Procedure Use the data presented in Figure 6.19 to plot ionic radius versus atomic number. Analyze and Conclude 1. Describe how the size changes when an atom forms a cation and when an atom forms an anion. 2. How do the ionic radii vary within a group of metals? How do they vary within a group of nonmetals? 3. Describe the shape of a portion of the graph that corresponds to one period. lonic radii (pm) 250 lonic Radii vs. Atomic Number 200 150 100 50 0 10 20 30 40 50 60 Atomic number 4. Is the trend across a period similar or different for periods 2, 3, 4, and 5? 5. Propose explanations for the trends you have described for ionic radii within groups and across periods. TA H 30 156 Li 2A BA 50 Figure 6.19 Atomic and ionic radii are an indication of the relative size of atoms and ions. The data listed in Figure 6.19 is reported in picometers (pm). He 3A 4A SA GA 7A 113 83 77 70 66 62 70 Be B C NO F Ne 14 2+ 3+ 60 44 23 15 146 140 133 191 160 143 109 109 105 102 Na Mg Al Si P S 94 CL Ar 156 Atomic radius Li 1+ 3+ 4+ 95 66 51 41 212 184 181 60 lonic radius 238 197 141 122 122 Ga Ge As 3+ 4+ 133 99 62 53 198 120 Se Br Kr 196 120 111 Metal 55 15 166 139 137 139 140 In Sn Sb Te I Xe Metalloid 130 Nonmetal 2+ 3+ 5+ 148 112 81 71 62 221 220 24 172 175 170 Ba TI Pb Bi 168 Po At Rn 140 140 Cation 24 3+ 4+ 5+ Anion 169 134 95 84 74 Section 6.3 Periodic Trends 175 Figure 6.20 This diagram Group 1A Group 7A compares the relative sizes of atoms and ions for selected alkali metals and halogens. The data are given in picometers. Comparing and Contrasting What happens to the radius when an atom forms a cation? Li F 156 60 62 133 Na+ CI CI When an atom forms an anion? 191 95 102 181 Figure 6.21 The ionic radii for cations and anions decrease from left to right across periods and increase from top to bottom within groups. 176 Chapter 6 Size generally increases 238 133 Br 120 196 Trends in lonic Size During reactions between metals and nonmetals, metal atoms tend to lose electrons and nonmetal atoms tend to gain electrons. The transfer has a predictable affect on the size of the ions that form. Cations are always smaller than the atoms from which they form. Anions are always larger than the atoms from which they form. Figure 6.20 compares the relative sizes of the atoms and ions for three metals in Group 1A. For each of these elements, the ion is much smaller than the atom. For example, the radius of a sodium ion (95 pm) is about half the radius of a sodium atom (191 pm). When a sodium atom loses an electron, the attraction between the remaining electrons and the nucleus is increased. The electrons are drawn closer to the nucleus. Also, metals that are representative elements tend to lose all their outermost electrons during ionization. So the ion has one fewer occupied energy level. The trend is the opposite for nonmetals like the halogens in Group 7A. For each of these elements, the ion is much larger than the atom. For exam- ple, the radius of a fluoride ion (133 pm) is more than twice the radius of a fluorine atom (62 pm). As the number of electrons increases, the attraction of the nucleus for any one electron decreases. Look back at Figure 6.19. From left to right across a period, two trends are visible-a gradual decrease in the size of the positive ions followed by a gradual decrease in the size of the negative ions. Figure 6.21 summarizes the group and periodic trends in ionic size. Trends in lonic Size Size of cations decreases Size of anions decreases Trends in Electronegativity In Chapters 7 and 8, you will study two types of bonds that can exist in compounds. Electrons are involved in both types of bonds. There is a prop- erty that can be used to predict the type of bond that will form during a reaction. This property is called electronegativity. Electronegativity is the ability of an atom of an element to attract electrons when the atom is in a compound. Scientists use factors such as ionization energy to calculate val- ues for electronegativity. Table 6.2 lists electronegativity values for representative elements in Groups 1A through 7A. The elements are arranged in the same order as in a periodic table. The noble gases are omitted because they do not form many compounds. The data in Table 6.2 is expressed in units called Paulings. Linus Pauling won a Nobel Prize in Chemistry for his work on chemical bonds. He was the first to define electronegativity. In general, electronegativity values decrease from top to bottom within a group. For representative elements, the values tend to increase from left to right across a period. Metals at the far left of the periodic table have low values. By contrast, nonmetals at the far right (excluding noble gases) have high values. The electronegativity values among the transition metals are not as regular. The least electronegative element is cesium, with an electronegativity value of 0.7. It has the least tendency to attract electrons. When it reacts, it tends to lose electrons and form positive ions. The most electronegative element is fluorine, with a value of 4.0. Because fluorine has such a strong tendency to attract electrons, when it is bonded to any other element it either attracts the shared electrons or forms a negative ion. ✔ chec Checkpoint) Why are values for noble gases omitted from Table 6.2? Table 6.2 Electronegativity Values for Selected Elements -Go nline SCLINKS ~CISTA SC For: Links on Electronegativity Visit: www.SciLinks.org Web Code: cdn-1063 H 2.1 Li Be B C 1.0 1.5 2.0 O N N 0 LL F 2.5 3.0 3.5 4.0 Na Mg AI Si P S CI 0.9 1.2 1.5 1.8 2.1 2.5 3.0 K Ca Ga Ge As Se Br 0.8 1.0 1.6 1.8 2.0 2.4 2.8 Rb Sr In Sn Sb Te 0.8 1.0 1.7 1.8 1.9 2.1 2.5 Cs Ba TI Pb Bi 0.7 0.9 1.8 1.9 1.9 Section 6.3 Periodic Trends 177 Figure 6.22 Properties that vary within groups and across periods include atomic size, ionic size, ionization energy, electronegativity, nuclear charge, and shielding effect. Interpreting Diagrams Which properties tend to decrease across a period? DO Atomic Mass com www size increases Ionic size increases lonization energy decreases Nuclear charge increases Electronegativity decreases Shielding increases Atomic size decreases lonization energy increases Electronegativity increases Nuclear charge increases Shielding is constant 1A 8A 2A ЗА 4A 5A 6A 7A Size of cations decreases Size of anions decreases Summary of Trends Figure 6.22 shows the trends for atomic size, ionization energy, ionic size, and electronegativity in Groups 1A through 8A. These properties vary within groups and across periods. The trends that exist among these proper- ties can be explained by variations in atomic structure. The increase in nuclear charge within groups and across periods explains many trends. Within groups an increase in shielding has a significant effect. 16. 17. 6.3 Section Assessment 18. 19. 20. 21. Key Concept How does atomic size change within groups and across periods? Key Concept When do ions form? Key Concept What happens to first ionization energy within groups and across periods? Key Concept Compare the size of ions to the size of the atoms from which they form. Key Concept How does electronegativity vary within groups and across periods? Key Concept In general, how can the periodic trends displayed by elements be explained? 22. Arrange these elements in order of decreasing atomic size: sulfur, chlorine, aluminum, and sodium. Does your arrangement demonstrate a periodic trend or a group trend? 23. Which element in each pair has the larger first ionization energy? a. sodium, potassium b. magnesium, phosphorus Writing Activity Explaining Trends in Atomic Size Explain why the size of an atom tends to increase from top to bottom within a group. Explain why the size of an atom tends to decrease from left to right across a period. Interactive Textbook Assessment 6.3 Test yourself on the concepts in Section 6.3. with ChemASAP 178 Chapter 6

6.1 Organizing the Elements 
Connecting to Your World 
In 1916, a self-service grocery store opened in Memphis, Tennessee. Shoppers could select items from shelves instead of waiting for a clerk to gather the items for them. In a self- 
service store, the customers must know how to find the products. From your experience, you know that products are grouped according to sim- ilar characteristics. You don't expect to find 
fresh fruit with canned fruit, or bottled 
juice with frozen juice. With a logical clas- sification system, finding and comparing products is easy. In this section, you will learn how elements are arranged in the periodic table and what that arrangement reveals about the elements. 
Searching For an Organizing Principle 
A few elements have been known for thousands of years, including copper, silver, and gold. Yet only 13 elements had been identified by the year 1700. Chemists suspected that other elements existed. They had even assigned names to some of these elements, but they were unable to isolate the ele- ments from their compounds. As chemists began to use scientific methods to search for elements, the rate of discovery increased. In one decade (1765-1775), chemists identified five new elements, including three color- less gases-hydrogen, nitrogen, and oxygen. Was there a limit to the num- ber of elements? How would chemists know when they had discovered all the elements? To begin to answer these questions, chemists needed to find a logical way to organize the elements. 
Chemists used the properties of elements to sort them into groups. In 1829, a German chemist, J. W. Dobereiner (1780-1849), published a clas- sification system. In his system, elements were grouped into triads. A triad is a set of three elements with similar properties. The elements in Figure 6.1 formed one triad. Chlorine, bromine, and iodine may look different. But they have very similar chemical properties. For example, they react easily with metals. Dobereiner noted a pattern in his triads. One element in each triad tended to have properties with values that fell midway between those of the other two elements. For example, the average of the atomic masses of chlorine and iodine is [(35.453 + 126.90)/2] or 81.177 amu. This value is close to the atomic mass of bromine, which is 79.904 amu. Unfortunately, all the known elements could not be grouped into triads. 
Guide for Reading 
Key Concepts 
PARMESANGASTEKKERS 
• How did chemists begin to 
organize the known elements? 
• How did Mendeleev organize his periodic table? 
• How is the modern periodic table organized? 
• What are three broad classes of elements? 
Vocabulary 
periodic law 
metals 
nonmetals 
metalloid 
Reading Strategy Comparing and Contrasting As you read, compare and contrast Figures 6.4 and 6.5. How are these two versions of the periodic table similar? How are they different? 
Figure 6.1 Chlorine, bromine, and iodine have very similar chemical properties. The numbers shown are the average atomic masses for these elements. 
Chlorine Bromine lodine 35.453 amu 79.904 amu 126.90 amu 
Section 6.1 Organizing the Elements 
155 
ГОЛЕТИЕ ПЕРИОДИЧЕСКОГО ЗАКОНА И.МЕНДЕЛЕЕВА 
Al=27.4 92=6869 Air & 116/13 
JOчTA CCCP 1969 OK 
Figure 6.2 Dimitri Mendeleev proposed a periodic table that could be used to predict the properties of undiscovered elements. 
Mendeleev's Periodic Table 
From 1829 to 1869, different systems were proposed, but none of them gained wide acceptance. In 1869, a Russian chemist and teacher, Dmitri Mendeleev, published a table of the elements. Later that year, a German chemist, Lothar Meyer, published a nearly identical table. Mendeleev was given more credit than Meyer because he published his table first and because he was better able to explain its usefulness. The stamp in Figure 6.2 is one of many ways that Mendeleev's work has been honored. 
Mendeleev developed his table while working on a textbook for his stu- dents. He needed a way to show the relationships among more than 60 ele- ments. He wrote the properties of each element on a separate note card. This approach allowed him to move the cards around until he found an organization that worked. The organization he chose was a periodic table. Elements in a periodic table are arranged into groups based on a set of repeating properties. Mendeleev arranged the elements in his periodic table in order of increasing atomic mass. 
Figure 6.3 is an early version of Mendeleev's periodic table. Look at the column that starts with Ti = 50. Notice the two question marks between the entries for zinc (Zn) and arsenic (As). Mendeleev left these spaces in his table because he knew that bromine belonged with chlorine and iodine. He predicted that elements would be discovered to fill those spaces, and he predicted what their properties would be based on their locations in the table. The elements between zinc and arsenic were gallium and germa- nium, which were discovered in 1875 and 1886, respectively. There was a close match between the predicted properties and the actual properties of these elements. This match helped convince scientists that Mendeleev's periodic table was a powerful tool. 
Figure 6.3 In this early version of Mendeleev's periodic table, the rows contain elements with similar properties. Observing A fourth element is grouped with chlorine (CI), bromine (Br), and (I) iodine. What is this element's symbol? 
156 Chapter 6 
но въ ней, мнѣ кажется, уже ясно выражается примѣнимость вы ставляемаго мною начала ко всей совокупности элементовъ, пай которыхъ извѣстенъ съ достовѣрностію. На этотъ разъ и и желалъ преимущественно найдти общую систему элементовъ. Вотъ этотъ 
OUNTE 
Ti-00 V=51 
Zr 90 Nb-94 
?=150. Ta-132. 
Mo=96 W=186. 
Mu-55 
Rh-1044 Pt-197,1 
Fe-56 
Ru=104, Ir-198. 
M-C6=59 
Pl=106, Os 199. 
Cu 631 
Be=9,1 B-11 
Mg-24 
Al=27,1 
Zu-65,2 
Ag=108 Hg=200. Cd=112 
2-68 
C-12 
Si=28 
Ur-116 Au=1972 S=118 
N=14 
P=31 
As=75 
Sb 122 
Bi 210 
0-16 
S-32 
F-19 
CI 35,5 
Li-7 Na 23 
K-39 
Se-79,4 Br 80 Rh-85,4 
Te=128? 
1-127 
Cs-133 TI=201 
Ca=40 
Sr 57,8 
Ba=137 Pb=207. 
2-45 
Ce 92 
?Er-56 
La-94 
7Y 60 
Di=95 
?In 75,6 
Th 118? 
1 H 
2 Li Be 
11 12 
3 Na Mg 
196 20 4 K Ca 
37 38 
5 Rb Sr 
Sc Ti 
23 24 25 26 27 
He 
10 
B C N O F Ne 
13: 14 
Al Si P 
16 17 18 
S Cl Ar 
28 29 30 31 32 33 34 35 36 
V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 
39 40 41 42 43 44 
42 43 44 45 46 47 48 49 50 51 52 53 64 Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 
72 
55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 
12 23 74 75 76 77 
78 79 n 
7 80 81 82 83 84 85 86 6 Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 
87 88 89 90 91 92 93 
108 109 110 111 112 
7 Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Uuu Uub 
Th Pa U Np Pu 95 96 97 98 99 100 101 102 103 104 105 106 107 1 
The Periodic Law 
The atomic mass of iodine (I) is 126.90. The atomic mass of tellurium (Te) is 127.60. Based on its chemical properties, iodine belongs in a group with bromine and chlorine. So Mendeleev broke his rule and placed tellurium before iodine in his periodic table. He assumed that the atomic masses for iodine and tellurium were incorrect, but they were not. Iodine has a smaller atomic mass than tellurium does. A similar problem occurred with other pairs of elements. The problem wasn't with the atomic masses but with using atomic mass to organize the periodic table. 
Mendeleev developed his table before scientists knew about the struc- ture of atoms. He didn't know that the atoms of each element contain a unique number of protons. Remember that the number of protons is the atomic number. In 1913, a British physicist, Henry Moseley, determined an atomic number for each known element. Tellurium's atomic number is 52 and iodine's is 53. So it makes sense for iodine to come after tellurium in the periodic table. In the modern periodic table, elements are arranged in order of increasing atomic number. 
The elements in Figure 6.4 are arranged in order of atomic number, starting with hydrogen, which has atomic number 1. There are seven rows, or periods, in the table. Period 1 has 2 elements, Period 2 has 8 elements, Period 4 has 18 elements, and Period 6 has 32 elements. Each period corre- sponds to a principal energy level. There are more elements in higher num- bered periods because there are more orbitals in higher energy levels. (Recall the rules you studied in Chapter 5 for how electrons fill orbitals.) 
The elements within a column, or group, in the periodic table have similar properties. The properties of the elements within a period change as you move across a period from left to right. However, the pattern of properties within a period repeats as you move from one period to the next. This pattern gives rise to the periodic law: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties. 
Checkpoint) How many periods are there in a periodic table? Chec 
114 
Uug 
Figure 6.4 In the modern periodic table, the elements are arranged in order of increasing atomic number. Interpreting Diagrams 
How many elements are there in the second period? 
Word Origins 
Periodic comes from the Greek roots peri meaning "around" and hodos, mean- ing "path." In a periodic 
table, properties repeat from left to right across each period. The Greek word 
metron means "measure." What does perimeter 
mean? 
Section 6.1 Organizing the Elements 157 
-≤ -1 
1 
IA 
1A 
18 
VIIB 
8A 
13 
14 
H 
2 
IIA 
2A 
IIIB 
IVB 
VB 
Metals 
Metalloids 
Nonmetals 
3A 
4A 
3 
4 
Li 
Be 
B C 
3 
4 
5 
6 
7 
8 
9 
10 
11 
11 
12 
HIA 
IVA 
VA 
VIA 
VIA 
VIIIA 
18 
22 BUDE 
Na Mg 
38 
4B 
5B 
6B 
7B 
8B 
19 
20 
21 
22 
23 
24 
25 
26 
27 
28 
K 
Ca 
Sc 
Ti 
V 
Cr 
Mn 
Fe 
Co 
Ni 
37 
38 
39 
40 
41 
42 
43 
44 
45 
46 
Rb 
Sr 
Y 
Zr 
Nb Mo 
Tc 
Ru 
Rh 
55 
56 
71 
72 
73 
74 
75 
78 
77 
78 
འ༅ལ མ་དང 
1B 
29 
2208 
12 
13 
14 
AI 
Si 
30 
31 
32 
Cu Zn Ga 
Ge 
47 
48 
49 
50 
Pd Ag Cd 
In 
Sn Sb 
79 
80 
81 
82 
Cs Ba Lu 
Hf 
Ta 
W 
Re 
Os 
Ir 
Pt 
87 
88 
103 
104 
105 
106 
107 
108 
109 
110 
111 
Au Hg TI 
112 
Pb 
C D G 2- 563 
15 
16 
17 
IVB 
VIB 
2 
SA 
6A 
7A 
He 
7. 
8 
10 
N 
F 
Ne 
15 
16 
17 
18 
p 
S 
CI 
CLA Ar 
33 
34 
35 
36 
As 
Se 
Br 
Kr 
51 
32 
54 
Te 
I 
Xe 
83 
84 
85 
86 
Bi 
Po 
At 
Rn 
Fr 
Ra 
Lr 
Rf 
Db Sg 
Bh 
Hs 
Mt 
Ds 
Uuu Uub 
114 
Uuq 
57 
58 
59 
60 
61 
62 
63 
64 
65 
66 
67 
68 
69 
70 
La 
Ce 
Pr 
Nd Pm Sm Eu Gd 
Tb Dy Ho 
Er Tm Yb 
89 
90 
91 
92 
93 
94 
95 
96 
97 
98 
99 
100 
101 
102 
Ac 
Th 
Pa 
U 
Cf 
Es 
Fm Md No 
Figure 6.5 One way to classify elements in the periodic table is as metals, nonmetals, and metalloids. Inferring What is the purpose for the black stair- step line? 
Go 
Online 
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Nonmetals 
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Np Pu Am Cm Bk 
Metals, Nonmetals, and Metalloids 
Most periodic tables are laid out like the one in Figure 6.5. Some elements from Periods 6 and 7 are placed beneath the table. This arrangement makes the periodic table more compact. It also reflects an underlying structure of the periodic table, which you will study in Section 6.2. Each group in the table in Figure 6.5 has three labels. Scientists in the United States used the labels shown in red. Scientists in Europe used the labels shown in blue. There is some overlap between the systems, but in many cases two differ- ent groups have the same letter and number combination. 
For scientists to communicate clearly, they need to agree on the stan- dards they will use. The International Union of Pure and Applied Chemis- try (IUPAC) is an organization that sets standards for chemistry. In 1985, IUPAC proposed a new system for labeling groups in the periodic table. They numbered the groups from left to right 1 through 18 (the black labels in Figure 6.5). The large periodic table in Figure 6.9 includes the IUPAC sys- tem and the system used in the United States. The latter system will be most useful when you study how compounds form in Chapters 7 and 8. 
→ 
Dividing the elements into groups is not the only way to classify them based on their properties. The elements can be grouped into three broad classes based on their general properties. Three classes of elements are metals, nonmetals, and metalloids. Across a period, the properties of ele- ments become less metallic and more nonmetallic. 
Metals The number of yellow squares in Figure 6.5 shows that most ele- ments are metals-about 80 percent. Metals are good conductors of heat and electric current. A freshly cleaned or cut surface of a metal will have a high luster, or sheen. The sheen is caused by the metal's ability to reflect light. All metals are solids at room temperature, except for mercury (Hg). Many metals are ductile, meaning that they can be drawn into wires. Most metals are malleable, meaning that they can be hammered into thin sheets without breaking. Figure 6.6 shows how the properties of metals can deter- mine how metals are used. 
158 Chapter 6 
Nonmetals In Figure 6.5, blue is used to identify the nonmetals. These elements are in the upper-right corner of the periodic table. There is a greater variation in physical properties among nonmetals than among metals. Most nonmetals are gases at room temperature, including the main components of air-nitrogen and oxygen. A few are solids, such as sulfur and phosphorus. One nonmetal, bromine, is a dark-red liquid. 
The variation among nonmetals makes it difficult to describe one set of general properties that will apply to all nonmetals. However, nonmetals are not metals, as their name implies. So they tend to have properties that are opposite to those of metals. In general, nonmetals are poor conductors of heat and electric current. Carbon is an exception to this rule. Solid nonmet- als tend to be brittle, meaning that they will shatter if hit with a hammer. 
Checkpoint Which type of elements tend to be good conductors of heat 
and electric current? 
Iron (Fe) 
The Gateway Arch in St. Louis, Missouri, is covered in stainless steel containing iron and two other metals, chromium (Cr) and nickel (Ni). The steel is shiny, malleable, and strong. It also resists rusting. 
Figure 6.6 The metals iron, copper, and aluminum have many important uses. How each metal is used is determined by its properties. 
Copper (Cu) 
Copper is ductile and second to only silver as a conductor of electric current. The copper used in electrical cables must be 99.99% pure. 
129.** 
Aluminum (AI) 
Aluminum is one of the metals that can be shaped into a thin sheet, or 
foil. To qualify as a foil, a metal must be no thicker than about 0.15 mm. 
Section 6.1 Organizing the Elements 159 
Figure 6.7 Pancake-sized circular slices of silicon, called wafers, are used to make computer chips. Because a tiny speck of dust can ruin a wafer, the people who handle the wafers must wear "bunny" suits. The suits prevent skin, hair, or lint from clothing from entering the room's atmosphere. 
Metalloids There is a heavy stair-step line in Figure 6.5 that separates the metals from the nonmetals. Most of the elements that border this line are shaded green. These elements are metalloids. A metalloid generally has properties that are similar to those of metals and nonmetals. Under some conditions, a metalloid may behave like a metal. Under other conditions, it may behave like a nonmetal. The behavior often can be controlled by changing the conditions. For example, pure silicon is a poor conductor of electric current, like most nonmetals. But if a small amount of boron is mixed with silicon, the mixture is a good conductor of electric current, like most metals. Silicon can be cut into wafers, like those being inspected in Figure 6.7, and used to make computer chips. 
1. 
2. 
3. 
4. 
6.1 Section Assessment 
Key Concept How did chemists begin the process of organizing elements? 
Key Concept What property did Mendeleev use to organize his periodic table? 
Key Concept How are elements arranged in the modern periodic table? 
Key Concept Name the three broad classes of elements. 
5. Which of these sets of elements have similar 
physical and chemical properties? 
a. oxygen, nitrogen, carbon, boron 
b. strontium, magnesium, calcium, beryllium c. nitrogen, neon, nickel, niobium 
6. Identify each element as a metal, metalloid, 
or nonmetal. 
a. gold 
c. sulfur 
160 Chapter 6 
b. silicon 
d. barium 
7. Name two elements that have properties similar 
to those of the element sodium. 
Connecting 
Concepts 
Atomic Number What does an atomic number tell you about the atoms of an element? Why is atomic number better than atomic mass for organizing the elements in a periodic table? Use what you learned in Section 4.2 to answer this question. 
Inter 
teractive Textbook 
Assessment 6.1 Test yourself on the concepts in Section 6.1. 
with ChemASAP 
6.2 Classifying the Elements 
Connecting to Your World 
The sculptor Augustus Saint- Gaudens designed this gold coin at the request of Theodore Roosevelt. President Roosevelt wanted coins minted in the United States to be as beautiful as ancient Greek coins, which he admired.The coin is an example of a double eagle. The name derives from the fact that the coin was worth twice as much as $10 
coins called eagles. A coin may contain a lot of information in a small space-its 
value, the year it was minted, and its 
country of origin. Each square in a peri- odic table also contains a lot of informa- 
tion. In this section, you will learn what types of information are usually listed in a periodic table. 
Squares in the Periodic Table 
The periodic table displays the symbols and names of the elements, along with information about the structure of their atoms. Figure 6.8 shows one square from the detailed periodic table of the elements in Figure 6.9 on page 162. In the center of the square is the symbol for sodium (Na). The atomic number for sodium (11) is above the symbol. The element name and average atomic mass are below the symbol. There is also a verti- cal column with the numbers 2, 8, and 1, which are the number of electrons in each occupied energy level of a sodium atom. 
The symbol for sodium is printed in black because sodium is a solid at room temperature. In Figure 6.9, the symbols for gases are in red. The sym- bols for the two elements that are liquids at room temperature, mercury and bromine, are in blue. The symbols for some elements in Figure 6.9 are printed in green. These elements are not found in nature. In Chapter 25, you will learn how scientists produce these elements. 
The background colors in the squares are used to distinguish groups of elements. For example, two shades of gold are used for the metals in Groups 1A and 2A. The Group 1A elements are called alkali metals, and the Group 2A elements are called alkaline earth metals. The name alkali comes from the Arabic al aqali, meaning "the ashes." Wood ashes are rich in com- pounds of the alkali metals sodium and potassium. Some groups of non- metals also have special names. The nonmetals of Group 7A are called halogens. The name halogen comes from the combination of the Greek word hals, meaning salt, and the Latin word genesis, meaning "to be born." There is a general class of compounds called salts, which include the com- pound called table salt. Chlorine, bromine and iodine, the most common halogens, can be prepared from their salts. 
Guide for Reading 
Key Concepts 
• What type of information can be displayed in a periodic table? 
• How can elements be classified based on their electron configurations? Vocabulary 
alkali metals 
alkaline earth metals halogens 
noble gases 
representative elements 
transition metal 
inner transition metal 
Reading Strategy Relating Text and Visuals 
As you read, look carefully at Figure 6,9. After you read the section, explain what you can tell about an element from the color assigned to its square and the color assigned to its symbol. 
Figure 6.8 This is the element square for sodium from the periodic table in Figure 6.9. Interpreting Diagrams What does the data in the square tell about the structure of sodium atoms? 
you 
11 
Na 
Sodium 22.990, 
-Atomic number 
Electrons in each energy level 
Element symbol 
-Element name 
Average atomic mass 
Section 6.2 Classifying the Elements 161 
1 
1A 
1 
H 
Hydrogen 1.0079 
1 
2 
2A 
2 
2 
N N 
Periodic Table of the Elements 
Representative Elements 
Transition Elements 
Alkali Metals 
Transition Metals 
C Solid 
Alkaline Earth Metals 
Inner transition metals 
Other Metals 
Br Liquid 
Hel Gas 
Metalloids 
Li 
Lithium 
6.941 
11 
Be 
Beryllium 9.0122 
12 
Na Mg 
Sodium 
22.990 
19 
K 
Potassium 39.098 
37 
Magnesium 24.305 
20 
Ca 
Calcium 
40.08 
Nonmetals 
Noble Gases 
Not found 
Te 
in nature 
2 
NJ 00 NJ 
8 
2 
3 
4 
5 
6 
7 
N 00 00 N 
8 
2 
24 0 0 24 
2 
3B 
21 
Sc 
Scandium 
22 
Ti 
Titanium 
44.9564790 
8 
10 
4B 
5B 
6B 
7B 
9 
8B 
8 
$ 
2002 
23 
V 
na 23.00 ho 
24 
25 
26 
27 
Vanadium 
50.941 
2 
2 
2 
38 8 
39 8 
40 
41 
8. 
42 
18 
19 
18 
18 
18 
18 
Rb 
8 
1 
Sr 
મૈં 
Y 
9 
Zr 
10 
2 
Nb 
12 
1 
Mo 
13 
17: 
Cr 
Chromium 51.996 
Manganese 54.938 
13 
15 
Mn2 
Fe 
Co 
NGON 
2 
Iron 
55.847 
Cobalt 
58.933 
2883. 
2 
43 
18. 
14. 
- $0.00 
44 
45 
Ru 
18 15 
18 
Rh 
16 
1 
Rubidium 
85.468 
55 
Cs 
Cesium 
132.91 
18 
Strontium 
87.62 
Yttrium 88.906 
2 
18 
100 00 00 00 - 
56 
8 
71 
Zirconium 
91.22 
272 
Niobium 
92.906 
સ 
8 
Molybdenum 95.94 
74 2 
Technetium (98) 
Ruthenium 
101:07: 
Rhodium 
102.91: 
75 ? 
76 
77 2 
18 
18 
18 
18 
8 
Ba 
18 
મ 
Lu 
32 
Hf 
10 
32 Ta 
32 
17 
W 
18 
18 
18 
18 
32 12 
Re 
32 
13 
Os 
32 
14 
Ir 
32 
15 
2 
2 
2. 
2 
2 
2 
NONGON 
8 
2 
Barium 
137.33 
Lutetium 
174.97 
Hafnium 
178.49 
Tantalum 180.95 
Tungsten 183.85 
Rhenium 186.21. 
Osmium 190.2 
Iridium 
192.22 
2 
87 
8 
88 8 
103 
104 
105 
106 
107 
108 
18 
18 
18 
18 
18 
18: 
Fr 
32 
18 
Ra 
32 
32. 
32 
18 
32 
32 
Db 
32 
32 
32 
32 
Sg Bh Hs 
32 
32 
8 
8 
9 
10 
Francium 1 
(223) 
Radium (226) 
2 
2 
Lawrencium Rutherfordium 
ን 
(262 
(261) 
Dubnium (262) 
11 2 
32 13 
32 
Seaborgium 2 
(263) 
Bohrium 2 
(264) 
14 Hassium 2 
(265) 
NANNOIN 
8 
109 
18 
Mt 
32 
32 
15 
Meitnerium 2 
200224 N 
8 
(268) 
Lanthanide Series 
57 
La 
58 
59 
60 
61 
62 
19 
18 
18 
Ce 
20 
Pr 
Nd 
Sm 24 
Lanthanum 138.91 
Cerium 140.12 
Praseodymium 140.91 
Neodymium 144.24 
Promethium (145) 
Samarium 150.4 
Actinide Series 
89 
90 
91 
92 
93 
94 
Ac 
Th 
19 32 
Pa 
321 20 
U 
32 
32 
24 
10 
Actinium 
N 
Thorium 
232.04 
Protactinium 
231 04 
Uranium 
238.03 
Neptunium (237) 
Plutonium 
(244) 
Figure 6.9 In this periodic table, the colors of the boxes are used to classify representative elements and transition elements. 
(227) 
162 Chapter 6 
Atomic number 
14 
Electrons in each 
energy level 
Si 
Element symbol 
13 
14 
15 
16 
17 
Silicon 
-Element name 
3A 
44 
5A 
6A 
7A 
18 8A 
2 
He 
Hehur 4.0028 
28.086 
5 
6 
7 
8 
9 
10 
Average 
atomic mass 
B 
C 
N 
0 
F 
Ne 
Boron 
10.81 
13 
Carbon 
12.011 
Nitrogen 
14.007 
Oxygen 15.999 
Fluorine 
18.998 
Neon 20.179 
200 
8 
14 
3 
ΑΙ 
Si 
P 
UG 
15 
evo in 
16 
6. 
10 
11 
12 
NOON. 
28 
Ni 
Nickel 
58.71 
46 
36 
18: 
18 
Pd 
18 
Palladium 
106.4 
78 
Pt 
Platinum 195.09 
1B 
29 
Cu 
Copper 63.546 
47 
18 
Zinc 
65.38 
48 
Ag Cd 
Silver 107.87 
79 
Cadmium 
112.41 
18 
18 
18 
Gallium 69.72 
49 
In 
Indium 
114.82 
81 
TI 
30 
18 
Zn 
ཀྑ ན བ ༦ 
2B 
Aluminum 
26.982 
Silicon 28.088 
Phosphorus 
30.974 
S 
Sulfur 32.06 
17 
CI 
Chlorine 
35.453 
Argon 
Ar 
2882 
2 
31 4 
32 
33 
18 
Ga 
3 
As 
04 00 00 00 102 
2 
2 
8 
51 
52 
53 
18 
18 
18 
18 
3 
Sn 
18 
Te 
I 
18 
Tin 
118.89 
Ge 
Germanium 72-59 
50 
*4 00 00 00 *4* 
Arsenic 74.922 
Sb 
Antimony 12175 
Selenium 
78.96 
35 
Br 
Bromine 79.904 
18 
28888 
36 
Kr 
Krypton 
lodine 
126.90 
Xe 
Xenon 13130 
34 
18 
Se 
70000 60 
2007 
Tellurium 
127.60 
80 
18 
18 
32 
Au 
32 
18 
Hg 
28832183 
2 
2 
2 
82 
B 
83 8 
84 8 
85 
36 
18 
18 
18 
Pb 
32 18 
Bi 
32 
18 
Po 
32 
18 
At 
Rn 
4 
$ 
6 
Gold 196.97 
Mercury 200.59 
Thallium 
Lead 
Bismuth 
204.37 
207.2 
208.98 
Polonium (209) 
Astatine (210) 
Radon 
110 8 
18 
32 
32 
NOONGE 
111 
32 
32 
18 
CONNO CON 
112 
114 
Uub 
18 
32 32 
18 
Darmstadtium 1 
(269) 
Unununium (272) 
Ununbium 
(277) 
*Uuq 
Ununquadium 
*Name not officially assigned. 
63 
Eu 
Europium 151.96 
Gadolinium 157.25 
64 
65 
66 
67 
63 
69 
70 
Gd 
25 
Tb 
Dy 
Ho 
Er 
Tm 
Yb 
Terbium 358.93 
Dysprosium 162.50 
Holmium 164.93 
Erbium 167.26 
Thulium 168.93 
Ytterbium 
173.04 
95 
96 
97 
98 
99 
100 
101 
102 
30 
Americium (243) 
Curium (247) 
Berkelium (247) 
Californium (251) 
Einsteinium (252) 
Fermium (257) 
Mendelevium 
(258) 
Nobelium 
12591 
Section 6.2 Classifying the Elements 163 
Figure 6.10 This blimp contains helium, one of the noble gases. Applying Concepts What does the ability of a helium- filled blimp to rise in air tell you about the density of helium? 
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Electron Configurations in Groups 
Electrons play a key role in determining the properties of elements. So there should be a connection between an element's electron configuration and its location in the periodic table. Elements can be sorted into noble gases, representative elements, transition metals, or inner transition metals based on their electron configurations. You may want to refer to Figure 6.9 as you read about these classes of elements. 
The Noble Gases The blimp in Figure 6.10 is filled with helium. Helium is an example of a noble gas. The noble gases are the clements in Group 9A of the periodic table. These nonmetals are sometimes called the inert gases because they rarely take part in a reaction. The electron configurations for the first four noble gases in Group 8A are listed below. 
Helium (He) 
Neon (Ne) 
Argon (Ar) 
Krypton (Kr) 
152 1s22s22p 1s22s22p 3s23p 1s22s22p63s23p63d1o4s24p6 
Look at the description of the highest occupied energy level for each ele- ment, which is highlighted in yellow. The s and p sublevels are completely filled with electrons. Chapter 7 will explain how this arrangement of elec- trons is related to the relative inactivity of the noble gases. 
The Representative Elements Figure 6.11 shows the portion of the periodic table containing Groups 1A through 7A. Elements in these groups are often referred to as representative elements because they display a wide range of physical and chemical properties. Some are metals, some are non- metals, and some are metalloids. Most of them are solids, but a few are gases at room temperature, and one, bromine, is a liquid. 
In atoms of representative elements, the s and p sublevels of the high- est occupied energy level are not filled. Look at the electron configurations for lithium, sodium, and potassium. In atoms of these Group 1A elements, there is only one electron in the highest occupied energy level. The electron is in an s sublevel. 
Lithium (Li) Sodium (Na) Potassium (K) 
152251 1522522p63s1 1s22s22p3s23p°45' 
In atoms of carbon, silicon, and germanium, in Group 4A, there are four electrons in the highest occupied energy level. 
Carbon (C) 
Silicon (Si) Germanium (Ge) 
1s22s22p2 1s22522p63523p2 
1s22s22p63523p 3d1o4s24p2 
For any representative element, its group number equals the number of electrons in the highest occupied energy level. 
Check 
Checkpoint) Why are noble gases sometimes referred to as inert gases? 
164 Chapter 6 
Magnesium This magnified view of a leaf shows the green structures where light energy is changed into chemical energy. The compound chlorophyll, which contains magnesium, absorbs the light. 
Sodium When salt lakes evaporate, they form salt pans like this one in Death. Valley, California. The main salt in a salt pan is sodium chloride. 
Figure 6.11 Some of the representative elements 
exist in nature as elements. 
Others are found only 
' in compounds. 
Arsenic This bright red. ore is a major source of arsenic in Earth's crust. It contains a compound of arsenic and sulfur. 
H 
Hydrogen 
BA 
8 
Li 
Be 
B 
C 
N 
9.0122 
- 12 
Lithium 4941 
11 
Beryllium 
Na Mg 
Godium 72960 
Magnesium 24.305 
Aluminum 26.062 
Porad 
1051 
Carbon 12.011 
Nitrogen 14,002 
Oxygen 15.599 
Florine 18.598 
13 
15 
16. 
17. 
AI 
Si 
P 
S 
CI 
Silicon Phosphorus 
Sillfor 
Chlorine: 
30.974 
32.06 
38.953 
* 
13 
.20 
31 
34 
35 
B 
Patessturn 
K 
Ca 
Calcium 40.08 
Gallium 60.72 
Ga 
Ge 
As 
Se 
Hin 
74922 
Selenium 78.06 
Bronine 79.204 
2 
38 
49 
50 » 
53 
Rb Sr 
In Sn 
Sb 
Te 
I 
Sariu 
137.33 
Rukidhan 
Strontium 87.62 
indium 114.82 
Tin 
118.69 
55 
56 
B1 
話 
82 
Cs 
Ba 
TI 
Pb 
Cesium 
Thallium 
Lead 
13201 
204,37 
207.2 
88 
Fr 
Ra 
Fransdam (223) 
Azdium (220) 
Anthony 
2175 
52200 
lodine 126.90 
* 1 * * 
83 
84 
85 
Bi 
73 
Po 
At 
* 
Bismuth 
208.92 
Poloniuni (209) 
Astatine 
2210 
Sulfur These scientists are sampling gases being released from a volcano through a vent called a fumarole. The yellow substance is sulfur. 
Transition Elements 
In the periodic table, the B groups separate the A groups on the left side of the table from the A groups on the right side. Elements in the B groups, which provide a connection between the two sets of representative ele- ments, are referred to as transition elements. There are two types of transi- tion elements-transition metals and inner transition metals. They are classified based on their electron configurations. 
The transition metals are the Group B elements that are usually dis- played in the main body of a periodic table. Copper, silver, gold, and iron are transition metals. In atoms of a transition metal, the highest occupied s sublevel and a nearby d sublevel contain electrons. These elements are characterized by the presence of electrons in d orbitals. 
The inner transition metals appear below the main body of the peri- odic table. In atoms of an inner transition metal, the highest occupied s sublevel and a nearby ƒ sublevel generally contain electrons. The inner transition metals are characterized by ƒ orbitals that contain electrons. Before scientists knew much about inner transition metals, people began to refer to them as rare-earth elements. This name is misleading because some inner transition metals are more abundant than other elements. 
Blocks of Elements If you consider both the electron configurations and the positions of the elements in the periodic table, another pattern emerges. In Figure 6.12, the periodic table is divided into sections, or blocks, that correspond to the highest occupied sublevels. The s block con- tains the elements in Groups 1A and 2A and the noble gas helium. The p block contains the elements in Groups 3A, 4A, 5A, 6A, 7A, and 8A, with the exception of helium. The transition metals belong to the d block, and the inner transition metals belong to the ƒ block. 
You can use Figure 6.12 to help determine electron configurations of elements. Each period on the periodic table corresponds to a principal energy level. Say an element is located in period 3. You know that the sand p sublevels in energy levels 1 and 2 are filled with electrons. You read across period 3 from left to right to complete the configuration. For transition ele- ments, electrons are added to a d sublevel with a principal energy level that is one less than the period number. For the inner transition metals, the principal energy level of the fsublevel is two less than the period number. This procedure gives the correct electron configurations for most atoms. 
Figure 6.12 This diagram classifies elements into blocks according to sublevels that are filled or filling with electrons. Interpreting Diagrams In the highest occupied energy level of a halogen atom, how many electrons are in the p sublevel? 
s block 
p block 
d1 d2 d3 d4 d5 do do do do dio 
d block 
fblock 
f2 f3 f4 f5 fô f7 få fŷ f10 f11 f12f13f14 
166 Chapter 6 
CONCEPTUAL PROBLEM 6.1 
Using Energy Sublevels to Write Electron Configurations Nitrogen is an element that organisms need to remain healthy. How- ever, most organisms cannot obtain nitrogen directly from air. A few organisms can convert elemental nitrogen into a form that can be absorbed by plant and animal cells. These include bacteria that live in lumps called nodules on the roots of legumes. The photograph shows the nodules on a bean plant. Use Figure 6.12 to write the elec- tron configuration for nitrogen (N), which has atomic number 7. 
Analyze Identify the relevant concepts. 
For all elements, the atomic number is equal to the total number of electrons. For a representa- tive element, the highest occupied energy level is the same as the number of the period in which the element is located. From the group in which the element is located, you can tell how many electrons are in this energy level. 
Practice Problems 
b. strontium 
8. Use Figure 6.9 and Figure 6.12 to write the elec- 
tron configurations of the following elements. a. carbon 
c. vanadium 
(Hint: Remember that the principal energy level number for elements in the d block is always one less than the period number.) 
Solve Apply concepts to this situation. Nitrogen is located in the second period of the periodic table and in the third group of the p block. Nitrogen has seven electrons. Based on Figure 6.12, the configuration for the two elec- trons in the first energy level is 1s2. The config- uration for the five electrons in the second energy level is 2s 22p3. 
9. List the symbols for all the elements whose 
electron configurations end as follows. Each n represents an energy level. 
a. ns2np1 
b. ns2nps 
c. ns2npond2(n+1)s2 
6.2 Section Assessment 
10. Key Concept What information can be 
included in a periodic table? 
11. Key Concept Into what four classes can 
elements be sorted based on their electron configurations? 
12. Why do the elements potassium and sodium have 
similar chemical properties? 
13. Classify each element as a representative element, 
transition metal, or noble gas. 
a. 1s22s22p°3s23p63d1o4s24p 
b. 1s22s22p63s23p®3d®4s2 
c. 1s22s22po3s23p2 
14. Which of the following elements are transition 
metals: Cu, Sr, Cd, Au, Al, Ge, Co? 
15. How many electrons are in the highest occupied 
energy level of a Group 5A element? 
Elements 
Handbook 
Noble Gases Look at the atomic properties of noble gases on page R36. Use what you know about the structure of atoms to explain why the color produced in a gas discharge tube is different for each gas. 
Interactive 
Textbook 
Assessment 6.2 Test yourself on the concepts in Section 6.2. 
with ChemASAP 
Section 6.2 Classifying the Elements 167 
Technology & Society 
True Colors 
Paint consists essentially of a pigment, a binder, and a liquid in which the other components are dissolved or dispersed. The liquid keeps the mixture thin enough to flow. The binder attaches the paint to the surface being painted, and the pigment determines the color. Pigments may be natural or manufactured. They may be inorganic or organic. The same pigment can be used in a water-based or oil-based paint. Comparing and Contrasting Describe at least three differences between the cave painting and the painting by Jacob Lawrence. 
Yellow ochre 
168 Chapter 6 
Natural pigments A prehistoric artist had a limited choice of colors- black from charcoal and red, brown, and yellow from oxides 
of iron in Earth's crust.These oxides (or ochre) pigments are often referred to as earth tones. 
Prehistoric art Around 14,000 years ago, 
an artist painted this bison on the ceiling of a cave in Spain. It is about two meters long. 
Charred wood is a 
source of charcoal, 
Red ochre 
Manganese Violet 
Red Iron Oxide 
Cadmium Orange 
M 
Cobalt Blue 
Cobalt Yellow 
Zinc White 
Chromium 
Oxide Green 
Manufactured pigments Alchemists (and then chemists) made pigments that don't exist in nature. They also made purer versions of natural pigments. Many of these pigments contain transition metals. 
From pigments to paint Artists mixed manufactured pigments with binders and solvents to make paint. Although premixed paints became available around 1800, some artists, including Jacob Lawrence, continued to mix their own paints. 
The Builders, 1974, by Jacob Lawrence 
M 
6.3 Periodic Trends 
Guide for Reading 
Key Concepts 
• What are the trends among the elements for atomic size? 
• How do ions form? 
• What are the trends among the elements for first ionization energy, ionic size, and electronegativity? 
• What is the underlying cause of periodic trends? 
Vocabulary 
atomic radius 
ion 
cation 
Connecting to Your World 
An atom doesn't have a sharply 
defined boundary. So the radius of an atom cannot be measured directly. There are ways to estimate the sizes of atoms. In one method, a solid is bombarded with X rays, and the paths of the X rays 
are recorded on film. Sodium chloride (table salt) produced the geometric pattern in the photograph.Such a pattern can be used to calculate the position of nuclei in a solid. The distances between nuclei in a solid 
are an indication of the size of the parti- cles in the solid. In this section, you will learn how properties such as atomic size are related to the location of elements in the 
periodic table. 
anion 
ionization energy electronegativity Reading Strategy 
Building Vocabulary After you read this section, explain the difference between a cation and an anion. 
Figure 6.13 This diagram lists the atomic radii of 
seven nonmetals. An atomic radius is half the distance between the nuclei of two atoms of the same element when the atoms are joined. 
Trends in Atomic Size 
Another way to think about atomic size is to look at the units that form when atoms of the same element are joined to one another. These units are called molecules. Figure 6.13 shows models of molecules (molecular models) for seven nonmetals. Because the atoms in each molecule are identical, the distance between the nuclei of these atoms can be used to estimate the size of the atoms. This size is expressed as an atomic radius. The atomic radius is one half of the distance between the nuclei of two atoms of the same element when the atoms are joined. 
The distances between atoms in a molecule are extremely small. So the atomic radius is often measured in picometers. Recall that there are one trillion, or 1012, picometers in a meter. The molecular model of iodine in Figure 6.13 is the largest. The distance between the nuclei in an iodine mol- ecule is 280 pm. Because the atomic radius is one half the distance between the nuclei, a value of 140 pm (280/2) is assigned as the radius of the iodine atom. In general, atomic size increases from top to bottom within a group and decreases from left to right across a period. 
Distance between nuclei 
Nucleus 
Hydrogen (H2) 30 pm 
Oxygen (O2) 
Nitrogen (N2) 
66 pm 
70 pm 
Atomic radius 
Fluorine (F2) 62 pm 
Chlorine (Cl2) 102 pm 
Bromine (Br2) 120 pm 
lodine (12) 140 pm 
170 Chapter 6 
Atomic radius (pm) 
Atomic Radius Versus Atomic Number 
300 
Period 4 
Period 5 
Cs 
Period 3 
Rb 
Period 2 
250 
200 
150 
Period 
Li 
Na 
100 
g 
50 
He 
Sc 
Cd 
Zn 
0 
10 
20 
30 
40 
50 
60 
Atomic number 
Group Trends in Atomic Size In the Figure 6.14 graph, atomic radius is plotted versus atomic number. Look at the data for the alkali metals. and noble gases. The atomic radius within these groups increases as the atomic number increases. This increase is an example of a trend. 
As the atomic number increases within a group, the charge on the nucleus increases and the number of occupied energy levels increases. These variables affect atomic size in opposite ways. The increase in positive charge draws electrons closer to the nucleus. The increase in the number of occupied orbitals shields electrons in the highest occupied energy level from the attraction of protons in the nucleus. The shielding effect is greater than the effect of the increase in nuclear charge. So the atomic size increases. 
Periodic Trends in Atomic Size Look again at Figure 6.14. In general, atomic size decreases across a period from left to right. Each element has one more proton and one more electron than the preceding element. Across a period, the electrons are added to the same principal energy level. The shielding effect is constant for all the elements in a period. The increas- ing nuclear charge pulls the electrons in the highest occupied energy level closer to the nucleus and the atomic size decreases. Figure 6.15 summa- rizes the group and period trends in atomic size. 
Trends in Atomic Size 
Size generally decreases 
Size generally increases 
Figure 6.14 This graph plots atomic radius versus atomic number for 55 elements. 
INTERPRETING GRAPHS 
a. Analyzing Data Which alkali metal has an atomic radius of 238 pm? 
b. Drawing Conclusions Based on the data for alkali metals and noble gases, how does atomic size change within a group? 
c. Predicting is an atom of barium, atomic number 56, smaller or larger than an atom of cesium (Cs)? 
Figure 6.15 The size of atoms tends to decrease from left to 
right across a period and increase 
from top to bottom within a group. Predicting If a halogen and an alkali metal are in the same period, which one will have the larger radius? 
Section 6.3 Periodic Trends 171 
Figure 6.16 When a sodium 
Nucleus 
11 p* 12 n° 
11 e 
Sodium atom (Na) 
atom loses an electron, it 
Nucleus 
becomes a positively charged 
17 p* 
ion. When a chlorine atom 
18 по 
17 
gains an electron, it becomes a negatively charged ion. Interpreting Diagrams What happens to the protons and neutrons during this change? 
lons 
Chlorine atom (CI) 
Lose one electron 
-1e- 
ر 
Gain one electron 
+1e 
10 e 
Nucleus 
+ 
11 pt 12 n° 
Sodium ion (Na') 
'Nucleus 17 p* 
18 @ 
18h0 
Chloride ion (CITM) 
teractive Textbook 
Animation 7 Discover the ways that atoms of elements combine to form compounds. 
with ChemASAP 
Some compounds are composed of particles called ions. An ion is an atom or group of atoms that has a positive or negative charge. An atom is electri- cally neutral because it has equal numbers of protons and electrons. For example, an atom of sodium (Na) has 11 positively charged protons and 11 negatively charged electrons. The net charge on a sodium atom is zero [(11+) + (−−11) = 0]. 
Positive and negative ions form when electrons are transferred between atoms. Atoms of metallic elements, such as sodium, tend to form ions by losing one or more electrons from their highest occupied energy levels. A sodium atom tends to lose one electron. Figure 6.16 compares the atomic structure of a sodium atom and a sodium ion. In the sodium ion, the number of electrons (10) is no longer equal to the number of protons (11). Because there are more positively charged protons than negatively charged electrons, the sodium ion has a net positive charge. An ion with a positive charge is called a cation. The charge for a cation is written as a number followed by a plus sign. If the charge is 1+, the number 1 is usually omitted from the complete symbol for the ion. So Na* is equivalent to Na1*. 
Atoms of nonmetallic elements, such as chlorine, tend to form ions by gaining one or more electrons. A chlorine atom tends to gain one electron. Figure 6.16 compares the atomic structure of a chlorine atom and a chlo- ride ion. In a chloride ion, the number of electrons (18) is no longer equal to the number of protons (17). Because there are more negatively charged electrons than positively charged protons, the chloride ion has a net nega- tive charge. An ion with a negative charge is called an anion. The charge for an anion is written as a number followed by a minus sign. 
Check 
Checkpoint) What is the difference between a cation and an anion? 
172 Chapter 6 
Trends in lonization Energy 
Recall that electrons can move to higher energy levels when atoms absorb energy. Sometimes there is enough energy to overcome the attraction of the protons in the nucleus. The energy required to remove an electron from an atom is called ionization energy. This energy is measured when an element is in its gaseous state. The energy required to remove the first electron from an atom is called the first ionization energy. The cation produced has a 1+ charge. First ionization energy tends to decrease from top to bottom within a group and increase from left to right across a period. 
Table 6.1 lists the first, second, and third ionization energies for the first 20 elements. The second ionization energy is the energy required to remove an electron from an ion with a 1+ charge. The ion produced has a 2+ charge. The third ionization energy is the energy required to remove an electron from an ion with a 2+ charge. The ion produced has a 3+ charge. 
Ionization energy can help you predict what ions elements will form. Look at the data in Table 6.1 for lithium (Li), sodium (Na), and potassium (K). The increase in energy between the first and second ionization ener- gies is large. It is relatively easy to remove one electron from a Group 1A metal atom, but it is difficult to remove a second electron. So Group 1A metals tend to form ions with a 1+ charge. 
Table 6.1 
Ionization Energies of First 20 Elements (kJ/mol) 
Symbol 
First 
Second 
Third 
H 
1312 
He (noble gas) 
2372 
5247 
Li 
520 
7297 
11,810 
Be 
899 
1757 
14,840 
B 
C 
801 
2430 
3659 
1086 
2352 
4619 
N 
1402 
2857 
4577 
1314 
3391 
5301 
F 
1681 
3375 
6045 
Ne (noble gas) 
2080 
3963 
6276 
Na 
496 
4565 
6912 
Mg 
738 
1450 
7732 
ΑΙ 
578 
1816 
2744 
Si 
786 
1577 
3229 
P 
1012 
1896 
2910 
S 
999 
2260 
3380 
K 
8 지곡으 
1256 
2297 
3850 
Ar (noble gas) 
1520 
2665 
3947 
419 
3069 
4600 
Ca 
590 
1146 
4941 
*An amount of matter equal to the atomic mass in grams. 
Section 6.3 Periodic Trends 
173 
Figure 6.17 This graph reveals. group and period trends for 
ionization energy. 
INTERPRETING GRAPHS 
a. Analyzing Data Which element in period 2 has the lowest first ionization energy? In period 3? 
b. Drawing Conclusions What is the group trend for first ionization energy for noble gases and alkali metals? c. Predicting If you drew a graph for second ionization energy, which element would you have to omit? Explain. 
terary.wwwAN, AARON She is gun wwww www.tw. EAT ADALA 
Figure 6.18 First ionization energy tends to increase from left to right across a period and decrease from top to bottom within a group. 
Predicting Which element would have the larger first ionization energy-an alkali metal in period 2 or an alkali metal in period 4? 
174 Chapter 6 
First ionization energy (kJ/mol) 
First lonization Energy Versus Atomic Number 
2500 
He 
Ne 
2000 
1500 
1000 
N 
Xe 
Be 
Zn As 
Cd 
Mg 
500 
Li 
Na 
K 
Rb 
CS 
0 
10 
20 
30 
40 
50 
60 
Atomic number 
Group Trends in lonization Energy Figure 6.17 is a graph of first ion- ization energy versus atomic number. Each red dot represents the data for one element. Look at the data for the noble gases and the alkali metals. In general, first ionization energy decreases from top to bottom within a group. Recall that the atomic size increases as the atomic number increases within a group. As the size of the atom increases, nuclear charge has a smaller effect on the electrons in the highest occupied energy level. So less energy is required to remove an electron from this energy level and the first ionization energy is lower. 
Periodic Trends in lonization Energy In general, the first ionization energy of representative elements tends to increase from left to right across a period. This trend can be explained by the nuclear charge, which increases, and the shielding effect, which remains constant. The nuclear charge increases across the period, but the shielding effect remains con- stant. So there is an increase in the attraction of the nucleus for an electron. Thus, it takes more energy to remove an electron from an atom. Figure 6.18 summarizes the group and period trends for first ionization energy. 
Trends in First lonization Energy 
Energy generally increases 
Energy generally decreases 
Quick LAB 
Periodic Trends in lonic Radii 
Purpose 
Make a graph of ionic radius versus atomic 
number and use the 
graph to identify periodic and group trends. 
Materials 
• graph paper 
Procedure 
Use the data presented in Figure 6.19 to plot ionic radius versus atomic number. 
Analyze and Conclude 
1. Describe how the size changes when an atom forms a cation and when an 
atom forms an anion. 
2. How do the ionic radii vary within a group of metals? How do they vary within a group of nonmetals? 
3. Describe the shape of a portion of the graph that corresponds to one period. 
lonic radii (pm) 
250 
lonic Radii vs. Atomic Number 
200 
150 
100 
50 
0 
10 20 30 40 50 60 
Atomic number 
4. Is the trend across a period similar or different for periods 2, 3, 4, and 5? 
5. Propose explanations for the trends you have described for ionic radii within groups and across periods. 
TA 
H 
30 
156 
Li 
2A 
BA 
50 
Figure 6.19 Atomic and ionic radii are an indication of the relative size of atoms 
and ions. The data listed in Figure 6.19 is reported in picometers (pm). 
He 
3A 
4A 
SA 
GA 
7A 
113 
83 
77 
70 
66 
62 
70 
Be 
B 
C 
NO 
F Ne 
14 
2+ 
3+ 
60 
44 
23 
15 
146 
140 
133 
191 
160 
143 
109 
109 
105 
102 
Na Mg 
Al 
Si P 
S 
94 
CL Ar 
156 
Atomic radius 
Li 
1+ 
3+ 
4+ 
95 
66 
51 
41 
212 
184 
181 
60 
lonic radius 
238 
197 
141 
122 
122 
Ga Ge 
As 
3+ 
4+ 
133 
99 
62 
53 
198 
120 
Se Br Kr 
196 
120 
111 
Metal 
55 
15 
166 
139 
137 
139 
140 
In Sn 
Sb 
Te 
I Xe 
Metalloid 
130 
Nonmetal 
2+ 
3+ 
5+ 
148 
112 
81 
71 
62 
221 
220 
24 
172 
175 
170 
Ba 
TI 
Pb Bi 
168 
Po At Rn 
140 
140 
Cation 
24 
3+ 
4+ 
5+ 
Anion 
169 
134 
95 
84 
74 
Section 6.3 Periodic Trends 
175 
Figure 6.20 This diagram 
Group 1A 
Group 7A 
compares the relative sizes of 
atoms and ions for selected alkali metals and halogens. The data are given in picometers. Comparing and Contrasting What happens to the radius 
when an atom forms a cation? 
Li 
F 
156 
60 
62 
133 
Na+ 
CI 
CI 
When an atom forms an anion? 
191 
95 
102 
181 
Figure 6.21 The ionic radii for 
cations and anions decrease from left to right across periods and increase from top to bottom within groups. 
176 Chapter 6 
Size generally increases 
238 
133 
Br 
120 
196 
Trends in lonic Size 
During reactions between metals and nonmetals, metal atoms tend to lose electrons and nonmetal atoms tend to gain electrons. The transfer has a predictable affect on the size of the ions that form. Cations are always smaller than the atoms from which they form. Anions are always larger than the atoms from which they form. 
Figure 6.20 compares the relative sizes of the atoms and ions for three metals in Group 1A. For each of these elements, the ion is much smaller than the atom. For example, the radius of a sodium ion (95 pm) is about half the radius of a sodium atom (191 pm). When a sodium atom loses an electron, the attraction between the remaining electrons and the nucleus is increased. The electrons are drawn closer to the nucleus. Also, metals that are representative elements tend to lose all their outermost electrons during ionization. So the ion has one fewer occupied energy level. 
The trend is the opposite for nonmetals like the halogens in Group 7A. For each of these elements, the ion is much larger than the atom. For exam- ple, the radius of a fluoride ion (133 pm) is more than twice the radius of a fluorine atom (62 pm). As the number of electrons increases, the attraction of the nucleus for any one electron decreases. 
Look back at Figure 6.19. From left to right across a period, two trends are visible-a gradual decrease in the size of the positive ions followed by a gradual decrease in the size of the negative ions. Figure 6.21 summarizes the group and periodic trends in ionic size. 
Trends in lonic Size 
Size of cations decreases 
Size of anions decreases 
Trends in Electronegativity 
In Chapters 7 and 8, you will study two types of bonds that can exist in compounds. Electrons are involved in both types of bonds. There is a prop- erty that can be used to predict the type of bond that will form during a reaction. This property is called electronegativity. Electronegativity is the ability of an atom of an element to attract electrons when the atom is in a compound. Scientists use factors such as ionization energy to calculate val- ues for electronegativity. 
Table 6.2 lists electronegativity values for representative elements in Groups 1A through 7A. The elements are arranged in the same order as in a periodic table. The noble gases are omitted because they do not form many compounds. The data in Table 6.2 is expressed in units called Paulings. Linus Pauling won a Nobel Prize in Chemistry for his work on chemical bonds. He was the first to define electronegativity. 
In general, electronegativity values decrease from top to bottom within a group. For representative elements, the values tend to increase from left to right across a period. Metals at the far left of the periodic table have low values. By contrast, nonmetals at the far right (excluding noble gases) have high values. The electronegativity values among the transition metals are not as regular. 
The least electronegative element is cesium, with an electronegativity value of 0.7. It has the least tendency to attract electrons. When it reacts, it tends to lose electrons and form positive ions. The most electronegative element is fluorine, with a value of 4.0. Because fluorine has such a strong tendency to attract electrons, when it is bonded to any other element it either attracts the shared electrons or forms a negative ion. 
✔ chec 
Checkpoint) Why are values for noble gases omitted from Table 6.2? 
Table 6.2 
Electronegativity Values for Selected Elements 
-Go 
nline 
SCLINKS 
~CISTA SC 
For: Links on 
Electronegativity 
Visit: www.SciLinks.org Web Code: cdn-1063 
H 
2.1 
Li 
Be 
B 
C 
1.0 
1.5 
2.0 
O N 
N 
0 
LL 
F 
2.5 
3.0 
3.5 
4.0 
Na 
Mg 
AI 
Si 
P 
S 
CI 
0.9 
1.2 
1.5 
1.8 
2.1 
2.5 
3.0 
K 
Ca 
Ga 
Ge 
As 
Se 
Br 
0.8 
1.0 
1.6 
1.8 
2.0 
2.4 
2.8 
Rb 
Sr 
In 
Sn 
Sb 
Te 
0.8 
1.0 
1.7 
1.8 
1.9 
2.1 
2.5 
Cs 
Ba 
TI 
Pb 
Bi 
0.7 
0.9 
1.8 
1.9 
1.9 
Section 6.3 Periodic Trends 
177 
Figure 6.22 Properties that vary within groups and across periods include atomic size, ionic size, ionization energy, electronegativity, nuclear charge, and shielding effect. Interpreting Diagrams Which properties tend to decrease across a period? 
DO Atomic 
Mass 
com 
www 
size increases 
Ionic size increases 
lonization energy decreases 
Nuclear charge increases Electronegativity decreases 
Shielding increases 
Atomic size decreases 
lonization energy increases Electronegativity increases 
Nuclear charge increases Shielding is constant 
1A 
8A 
2A 
ЗА 
4A 5A 6A 
7A 
Size of cations decreases 
Size of anions decreases 
Summary of Trends 
Figure 6.22 shows the trends for atomic size, ionization energy, ionic size, and electronegativity in Groups 1A through 8A. These properties vary within groups and across periods. The trends that exist among these proper- ties can be explained by variations in atomic structure. The increase in nuclear charge within groups and across periods explains many trends. Within groups an increase in shielding has a significant effect. 
16. 
17. 
6.3 Section Assessment 
18. 
19. 
20. 
21. 
Key Concept How does atomic size change within groups and across periods? 
Key Concept When do ions form? 
Key Concept What happens to first ionization energy within groups and across periods? 
Key Concept Compare the size of ions to the size of the atoms from which they form. 
Key Concept How does electronegativity vary within groups and across periods? 
Key Concept In general, how can the periodic trends displayed by elements be explained? 22. Arrange these elements in order of decreasing 
atomic size: sulfur, chlorine, aluminum, and sodium. Does your arrangement demonstrate a periodic trend or a group trend? 
23. Which element in each pair has the larger first 
ionization energy? 
a. sodium, potassium b. magnesium, phosphorus 
Writing 
Activity 
Explaining Trends in Atomic Size Explain why the size of an atom tends to increase from top to bottom within a group. Explain why the size of an atom tends to decrease from left to right across a period. 
Interactive 
Textbook 
Assessment 6.3 Test yourself on the concepts in Section 6.3. 
with ChemASAP 
178 Chapter 6

Transcript for:Understanding the Periodic Table and Trends

Transcript for:
Understanding the Periodic Table and Trends