welcome to this video about properties of the elements in this video we'll go over both chemical and physical properties of different elements and try to relate them to their positions on the periodic table first we will go over what differentiates a metal non-metal and a metaloid coent and ionic bonding and after that we will talk about the elements in groups 1 and 17 to give us an idea of how properties change down a group and then we'll talk about the elements across period 3 to give us an idea of how properties change across a period and finally we will end with a discussion on the acidic and basic oxides across periods 2 and three so let's jump straight into it the periodic table can be divided into Metals semi Metals otherwise known as metalloids and non-metals we will first first talk about the metals in metals atoms donate their outer electrons to a delocalized sea of electrons in metallic bonding the positive metal nuclei are mutually attracted to the bonding electrons metallic bonding thus requires the removal of electrons from the outer shells so it can only occur if the electro negativity and the ionization energy are small both parameters increase across a period and decrease down a group therefore metals are found to the left and bottom of the periodic table on the other hand non-metal atoms gain full outer shells by sharing electrons in directional coal bonds non-metal atoms contain many electrons in their outer shells so they can gain full outer shells by forming a small number of bonds an element is a non-metal if its electro negativity and ionization energy are high non-metals are only found at the top and right of the periodic table and what about the semimetals semimetals or metalloids are elements with properties between metals and non-metals they are often shiny like Metals but resemble non-metals in their chemical reactions they are also semiconductors becoming electrically conductive when the energy input or voltage is sufficiently high the metalloids are found in a zigzag from Boron in Period 2 to polonium and atine in Period 5 some groups contain a mixture of non-metals metals and metalloids for example group 14 contains a non-metal carbon two metalloids silicon and germanium and two metals tin and lead the metallic Behavior increases down this group now let's revise a couple more types of bonding bonding depends on the properties of the elements involved in that Bond the type of bonding between atoms depends on the electr negativity difference given the abbreviation Delta X ionic bonds occur if the electro negativity difference is large if the electro negativity difference Delta X is greater than 1.8 this is indicative of an ionic bond on the other hand Cove valent bonding occurs when this difference is small a Delta X of less than 1.8 is indicative of Cove valent bonding if this Delta X is between 0.5 and 1.8 the bonding electron density is much higher around one atom than the other making the bond Cove valent but polar for some metals the electro negativity is high enough for the atoms to form polar calent bonds this is true for the dblock metals and The Period 2 Metals lithium and burum for example burum chloride is calent even though magnesium chloride is ionic in this next section we will talk about a couple of the groups found on the periodic table we'll start with group one it contains the alkaline metals the alkal metals are soft Metals with low melting points the metals are in the same group so have similar chemical properties All Metals in group one lose one electron in reactions to form one plus ions for example sodium reacts with water to produce a sodium ion and and hydroxide ions and hydrogen gas according to this equation here the atomic radius increases down the group as the number of shells increases so for example lithium is smaller than sodium and sodium is smaller than potassium as the atomic radius increases the ionization energy decreases as the outermost electron becomes easier to remove as it's further away from the nucleus the decrease in ionization energy means less energy is needed to form the 1l ion in reactions so more energy is released overall this means that the metals become more reactive down the group whereas potassium reacts violently with water producing a lilac Flame the reaction with sodium may not be exothermic enough for the hydrogen to ignite and and what about physical properties alkal metals are soft and have low melting points because the atoms are large and each atom only contributes one electron for metallic bonding meaning that the attraction between the atoms and the bonding electrons are weak the number of outer electrons stays the same down the group but the atom size increases and electr negativity decreases therefore the delocalized electrons in the metallic bond are attracted less strongly and the melting point decreases on the other side of the periodic table we have group 17 which contains the halogens the halogens are non-metals with low melting and boiling points all halogens are normally found as datomic molecules hogen atoms have seven electrons in their outer shell so so forming a Cove valent Bond gives them a stable octet of electrons as in group one the atomic radius increases down the group and the electro negativity decreases bonding electrons in a hallogen halogen Bond are further from the nucleus when the hallogen atom is large and this results in a weaker and longer Bond the Florine Florine Bond however is an exception here the atoms are so small that there is a strong repulsion between the lone pairs of the two atoms causing this bond to be weaker than the chlorine chlorine bond the electron affinity becomes smaller I.E less negative down the group as less energy is released when an electron is added to the outer shell of the larger atoms therefore the elements near the top of the group form annion more readily this trend in electron affinity results in a reactivity series for reduction reactions the reactivity decreases down the group electrons are transferred to convert the most reactive hogen to its annion for example reacting Florine with chloride ions produces chlorine and fluoride ions according to this reaction equation here the more reactive h also react more readily with metals to form salts and what about the physical properties Florine and chlorine are gases bromine is a liquid and iodine is a solid the boiling point increases down the group even though the atoms are becoming larger unlike the alkali metals the halogens form simple Cove valent molecules that are help together by dispersion forces these forces are stronger for larger atoms as larger atoms have more electrons it's easier for these electrons to be redistributed within the atoms to produce induced partial charges in this section we will focus on period 3 as by looking at period 3 we can see how the properties of elements changes as we go across a period across a period the nuclear charge increases while the number of shells stays the same the outer shell electrons are therefore more strongly attracted to the nucleus causing the atomic radius to decrease this increased attraction between the veence electrons and the nucleus causes the ionization energy to increase and the electron affinity also becomes larger the the elements to the left of period 3 sodium to aluminum are metals while those to the right phosphorus to argon are non-metals silicon is a semimetal otherwise known as a metaloid the melting and boiling points of the elements are shown the metal atoms interact via metallic bonds the latice of positively charged nuclei are attracted to a sea of delocalized electrons melting point and hardness increase from sodium to aluminum because each atom contributes more veence electrons for metallic bonding and the atoms also become smaller so the nuclei are closer to the bonding electrons silicon atoms form a giant lce of directional coent bonds this causes silicon to be more brittle than the metals as it's more difficult for layers of atoms to slide over each other without breaking bonds the bonds are stronger than metallic bonds however so the melting point is much higher the non-metals in Period 3 all have lower melting points because with the exception of argon the atoms do not form a giant Cove valent lattice but instead form discrete molecules the atoms are linked by strong calent bonds but the molecules interact only by weak dispersion forces therefore very little energy is needed to overcome these interactions and separate the molecules phosphorus forms a number of allotropes in which atoms are linked in different ways the standard state is white phosphorus in which each molecule contains four atoms by contrast sulfur atoms form rings of eight compounds because these sulfur molecules are larger they form stronger dispersion forces and display a higher melting point chlorine is a smaller molecule still the molecule only contains two atoms linked by a single calent Bond the dispersion forces between the molecules are therefore very weak such that the molecule is a gas at room temperature Argan atoms do not readily form calent bonds because the atoms already have a full outer shell of electrons the element therefore consists of a single atom giving it an even lower melting point than chlorine in this last section we will discuss different oxides all elements in periods 2 and three with the exception of neon and argon as these are noble gases can react with oxygen sodium forms the oxide salt sodium oxide sodium peroxide and sodium super oxide which are are other oxides are also possible but much less stable magnesium reacts with oxygen to form the salt magnesium oxide these oxide salts react with water to form hydroxide ions in solution making them basic aluminum oxide is a giant Cove valent lattice with highly polar bonds because it's not an ionic solid it's insoluble in water aluminum oxide is OT IC this means that it acts as both an acid and a base reacting with both acids such as hot dilute aquous hydrochloric acid and bases such as hot concentrated aquous sodium hydroxide the oxides become less basic across a period because the elements become more electr negative and their bonds with oxygen become less ionic and more coal as electron density is drawn away from the oxygen ion or atom it's less able to donate electron density to a hydrogen ion in an acid base reaction the pH of a solution is a measure of the hydrogen ion concentration a small concentration gives a high pH basic oxides produce Solutions with a high pH because they react with water to produce hydroxide ions these hydroxide ions react with hydrogen ions to increase the pH non-metal oxides include NO2 p410 S2 and S SO3 nitrogen dioxide NO2 dissolves in water to form a mixture of nitrous acid and nitric acid the phosphorus oxide p410 is a simple calent compound with a cage likee structure similar to Diamond but one one that doesn't repeat in all directions it readily reacts with water to produce phosphoric acid a moderately strong acid sulfur reacts with oxygen to form sulfur dioxide and sulfur trioxide both are acidic gases they react with water to produce the weak acid sulfurous acid and the strong acid sulfuric acid the oxides become more iic across a period and down a group especially if you only consider the reactions of the most oxidized compounds for example the acid strengths of the oxides we mentioned increases in this order and that's it for this video on properties of the elements we'll quickly recap what we went over we learned where the metals non-metals and metalloids are on the periodic table and their different properties we also learned that ionic bonding happens when the electr negativity difference is high and calent bonding occurs when this difference is low we then discussed how properties change going down a group group one was representative of metallic groups and group 17 was representative of non-metallic groups next we talked about how properties change when going across a period using period perod 3 as an example we also learned that across a period the oxide character changes from basic to acidic so hopefully you learn something new about the properties of the elements in this video