Transcript for:
Chemistry Principles and Concepts Overview

Chemistry Section A: Principles of Chemistry 1) States of Matter 1.1) Matter is anything that has mass and occupies space. Matter can be solid, liquid or gaseous, depending on its physical properties. Physical properties are characteristics that can be observed without changing the composition of matter. e.g. color, odour, taste, etc. The Particulokinetic Theory of Matter states that: * Matter is made up of particles. * The particles are in constant random motion and tend to move away from each other. * A balance of forces between repulsion and attraction keep the particles together. Evidence for Particulokinetic Theory of Matter * Crystals have a regular shape. * Crystals dissolve. * Brownian Motion * Diffusion * Osmosis Diffusion Diffusion is the movement of particles from an area of high concentration to an area of low concentration. It results in a substance becoming evenly spread out. (Experiment 1.1, 1.2, 1.3) ________________ The experiment demonstrates that the less dense gas, ammonia, diffused faster than the more dense gas, hydrogen chloride. The white fumes are formed as a result of reaction between gaseous ammonia particles and gaseous hydrogen chloride particles. NH3 (g) + HCl (g) ——> NH4Cl (s) ammonia + hydrogen chloride ——-> ammonium chloride Brownian Motion Brownian motion is the haphazard motion of particles due to the bombardment by other invisible particles. The phenomenon takes its name from the British botanist Robert Brown, who noticed the random but vigorous motion of pollen grains while examining them in water using a microscope. It was explained by Albert Einstein. (Experiment 1.4) Osmosis Osmosis is the movement of solvent particles, usually water, from an area of high concentration to an area of low concentration through a semi-permeable membrane. A semi-permeable membrane allows small particles to pass through it, but not larger ones. Osmosis can be thought as diffusion in one direction. Can be used to explain the use of salt or sugar to control garden pests and as a preservative. (Experiment 1.5, 1.6) ________________ 1.2) Properties of Different States States of matter are never fixed and are determined by conditions of temperature and pressure which can affect: * the arrangement of the particles * the types of forces between the particles * the types of motion of which the particles are capable Physical State Property Explanation Solids Shape: * definite shape * many have three-dimensional lattices (i.e. they are crystalline) Volume: * constant volume Compressibility: * not easily compressible - pressure has no effect on volume * the arrangement is close-packed and regular * particles are strongly attracted * their motion is restricted * therefore the solid has a fixed shape and volume and is not easily compressible Liquids Shape: * variable shape - take the shape of their containers * no fixed lattice Volume: * constant volume Compressibility: * they are compressible to a small extent - pressure has small effect on volume * looser, irregular arrangement of particles * weaker inter-particle attractions * particles can move away from each other, either individually or as groups or clusters, dependent on how much energy they have * the shape of the liquid is therefore not fixed, but the volume is definite and liquids are not easily compressible because there is still significant attraction between the particles Gases Shape: * variable shape - take the shape of their containers Volume: * fill any space available Compressibility: * easily compressible * particles in a gas are ‘free agents’ * there is so little attraction between them that their energy is high and they can move freely, bouncing into each other and the walls of the container * gases therefore have no definite shape or volume and spread out to fill any available space * the particles in a gas are so far apart, they can be pushed closer together (i.e. volume decreases) when pressure is applied or fly apart (i.e. volume increases) when pressure is reduced * gases therefore are easily compressible 1.3) Most substances can exist in more than one state, depending on the conditions of temperature and pressure. Melting Particles vibrate more vigorously when a solid is heated. Eventually the vibrations become so violent that the forces of attraction no longer hold particles in position and the solid melts. The heat energy supplied during melting is used to break up the solid’s structure. The melting point of a substance is the temperature at which it changes from the solid state to the liquid state. The temperature remains constant during melting. Evaporation Evaporation is the process in which the particles of a liquid leave the surface of the liquid as a vapor. When a liquid is heated, the particles move faster, and continually collide with other particles. Occasionally, some of the particles acquire sufficient energy to break free of the surface and they escape. Escaping particles take with them a lot of energy, so evaporation leads to cooling. The rate of evaporation depends on: * the nature of the liquid * temperature * the amount of exposed surface Boiling Boiling is the process by which a liquid is freely converted to gas or vapor at its boiling point. When a pure liquid is heated its temperature rises until the boiling point is reached. Once boiling has started, the temperature remains steady. The heat energy supplied at the boiling point goes to separate the particles from each other so that they can enter the gas state. None of the energy supplied is used to raise the temperature of the liquid any further. Whereas evaporation occurs at the surface, boiling takes place throughout the liquid. Also, evaporation occurs spontaneously at all temperatures, but boiling occurs at one particular temperature for a given external pressure. ________________ Condensation The change from gas to liquid can be achieved by cooling for some gases (e.g. water vapour), or by a combination of cooling and compression for other gases (such as propane). As the gas cools, the particles lose energy and move closer together, thereby increasing the attraction between them. The effect of compression is that the increased pressure pushes the particles closer together. Freezing As a liquid cools, its particles lose energy and move closer together. This movement increases the attraction between the particles and allows for a more regular arrangement of its particles. That is when the liquid changes to a solid. Once the liquid begins to change to a solid, the temperature remains constant until all the liquid is converted. The temperature at which the solid and liquid are in equilibrium with each other at atmospheric pressure is known as the freezing point. Sublimation Sublimation is the change directly from a solid to a gas or vice versa, without going through the liquid state. 2) Mixtures and Separations 2.1) A pure substance is one that contains only one substance, either a single compound or a single element. Some characteristics of pure substances are: * definite and constant composition * definite and constant physical properties, such as fixed melting and boiling points, under a given set of conditions * distinct chemical properties * show only a single spot when analysed by chromatography A mixture is a material that contains two or more substances that are not chemically combined, such as mixtures of elements, of elements and compounds and of several compounds. Mixtures melt and boil over a range of temperatures rather than at fixed points. This is in keeping with their composition. The properties of the mixture are a combination of the properties of the different substances it contains. There are two broad categories of mixtures: * Homogeneous mixtures have the same composition throughout. e.g. a solution of sodium chloride in water, brass (an alloy of copper and zinc). A solution is a homogeneous mixture containing one or more solutes dissolved in a solvent. An alloy is a solution of a metal and another solid element. * Heterogeneous mixtures have a composition that is not uniform. e.g. a mixture of sand and water, milk (a mixture of butter fat and water). (Experiment 5.1, 5.2) 2.2) Solutions Any solution has two components: the solute and the solvent. The distinguishing features of solutions are: * solute and solvent are thoroughly mixed - all parts of the solution have the same chemical composition, chemical properties and physical properties * solute and solvent do not separate when the solution is allowed to stand * the particles of solute are not visible, even under an optical microscope * the solution may be coloured, but is usually transparent if the solvent is a liquid * the solute may, in many cases, be separated from the solvent by purely physical means Suspensions Examples of suspensions include mud in water, powdered chalk in water and some medicines that separate out into layers on standing. The distinguishing features of a suspension are: * the components separate out when the suspension is allowed to stand * the suspended particles are larger than those of solutions and colloids and are visible to the naked eye * the particles are not individual atoms or molecules, but are believed to be clusters of them Colloids Colloids can be thought of as being intermediate between true solutions and suspensions, as the particles of one component do not dissolve in the other and the particles do not separate out on standing. Examples of colloids are a starch/water mixture, toothpaste, smoke, whipped cream, soap suds and milk. The distinguishing features of colloids are: * the components do not separate out on standing * the components are not separated by simple filtration * the dispersed particles are intermediate in size between those of a solution and those of a suspension * the particles are clusters of molecules or atoms that are big enough to scatter a beam of light but too small to settle COMPARING SOLUTIONS, SUSPENSIONS, COLLOIDS Comparative Size of Particles (nm) Able to pass through filter paper? Settles out? Exhibits Brownian motion? Appearance Solutions 0.1-2 yes no yes transparent (if solvent is liquid); transparent Colloids 2-1000 yes not usually yes not transparent; may appear homogeneous; scatters Suspensions ≥1000 no yes no not transparent heterogeneous 2.3) TYPES OF SOLUTIONS Solute Solvent Examples gas liquid oxygen in water, carbon dioxide in fizzy drinks liquid liquid alcoholic drinks (alcohol in water), gasoline solid liquid sugar in water, iodine in ethanol solid solid alloys (brass, bronze, coinage metals, etc) ________________ 2.4) Solubility A solute is considered soluble if it dissolves readily in a given solvent. Sparingly soluble substances dissolve only to a small extent in a given solvent. The extent to which a solute dissolves in a particular solvent depends on: * the nature of the solute and the solvent * the temperature * the pressure The solubility of a solute is the number of grams of it that dissolves in 100g of solvent at a given temperature and pressure. It is generally the case that more solid dissolves as the temperature increases. However, the solubility of gases in liquids generally decreases as the temperature increases. Most gases increase their solubility as pressure increases. Graphs that show how solubility varies with temperature are known as solubility curves. The following procedures will increase the rate at which a solid dissolves in a solvent: * crushing - increases the surface of the solute exposed to the solvent * stirring - brings more solvent in contact with solute * heating - increases the movement of the solute particles, causing more mixing Concentration of Solutions The concentration of a solution is the amount of solute that is dissolved in a fixed volume of solution. A dilute solution contains a small quantity of solute dissolved in the solvent. A concentrated solution contains relatively large quantities of solute dissolved in the solvent. A saturated solution contains as much solute as the solvent can possibly dissolve at a particular temperature and pressure in the presence of undissolved solute. A supersaturated solution contains more solute than the solvent can normally dissolve at a given temperature and pressure. (Experiment 5.3) 2.5) Separation Techniques Filtration Filtration is a method used to separate suspended solids from a liquid. It is based on differences in particle size of the components. The filter paper acts as a selective physical barrier that allows the liquid to pass through it, but does not allow the solid to pass through. Filtration can be used if you want to keep either the liquid or the solid or both after the suspension. Sublimation Sublimation is used to separate a solid which sublimes from a mixture of solids. A mixture of sodium chloride and ammonium chloride can be separated by the process of sublimation. Ammonium chloride sublimes but sodium chloride does not. When a mixture of the two compounds is gently heated, the ammonium chloride collects on the base of the test tube whereas sodium chloride, which is unaffected by the gentle heat, remains behind. ________________ Use of Solvents Solids may also be separated from one another by the use of solvents. A solvent that dissolves one of the components of the mixture but not the other is chosen. In order to remove copper (II) sulphate from a mixture of sand and copper (II) sulphate: 1. Solution: water is added to dissolve the copper (II) sulphate. 2. Filtration: the sand is retrieved by filtration. 3. Crystallisation: the filtrate in this case is a solution of copper sulphate in water. It is then heated in an evaporating dish to concentrate the solution until it becomes saturated. The saturated solution is allowed to cool. As it does so, crystals separate out. Simple Distillation Simple distillation is used to obtain a pure solvent from a solution (e.g. the purification of water from sea water). The method depends on the fact that the solvent vaporizes at a much lower temperature than the solute. evaporation of solvent condensation solution —————> pure solvent (as vapour) ———————-> pure solvent (as liquid) heating cooling The solid remains in the distillation flask and is called the residue. The distilled solvent is the desired product and is called the distillate. Note: * condensation and boiling take place in different parts of the apparatus * the thermometer is in contact with vapour - in this position, the thermometer reads the temperature of the pure vapour * it is advisable to add anti-bumping granules to the flask to achieve steady boiling * water to cool the vapour enters the condenser from the end closer to the receiver, and leaves from the end close to the distilling flask. Fractional Distillation Fractional distillation is used to separate miscible liquid where the components of the liquid mixture have boiling points that are close together (e.g. ethanol and water). It can be used on an industrial scale to separate the components of liquid air as well as in the refining of crude oil. Note: * the vapour passes through a column of glass beads (called a fractionating column) before going through the condenser * the temperature of the fractionating column decreases as you go from bottom to top * the thermometer is in contact with the vapour as it enters the condenser ________________ Process: * as the liquid mixture is heated, the components vaporise * ethanol has a lower boiling point and is therefore more volatile * the vapour contains more of the more volatile component (ethanol) * the mixture of vapours passes up the fractionating column * the vapour condenses and vaporises in the column many times as it rises * the more volatile component eventually rises to the top of the column as vapor while the less volatile component condenses and falls down the column * the more volatile component passes through the condenser and is collected as liquid in the receiving flask Using a Separating Funnel A separating funnel is used to separate liquids that are immiscible. The method is based on the fact that immiscible liquids do not mix but form two distinct layers, with the less dense liquid on top and the more dense liquid below. Note: * the mixture of immiscible liquids is placed in the funnel * the components are left to separate * the denser liquid is withdrawn through the tap into one container * the tap is closed just before the second layer gets to it * the container is substituted and a little liquid is allowed to run out until only the upper layer begins to flow * the container is changed again and the upper liquid is run out and collected Solvent Extraction Solvent extraction is used to separate a component from a mixture by using two solvents. Solvent extraction depends on: * the desired component being more soluble in one solvent than in the other * the two solvents being immiscible Solvent extraction is a good method to extract organic substances from aqueous solutions (e.g. removal of caffeine from tea or coffee). Caffeine is a stimulant that affects the nervous system. It is an organic compound that is more soluble in dichloromethane than it is in water. (Experiment 5.4, 5.5) Paper Chromatography Paper chromatography is a technique used to separate, purify or identify substances. This method involves the use of a stationary phase and a mobile phase. The mixture to be separated is placed on a strip of filter paper (or chromatography paper) and a solvent is allowed to move through it. The separation is based on the differences in the rates of movement of the different components in the mixture along the paper. To separate substances by paper chromatography, a fine capillary tube is used to place a spot of the solution containing the mixture to be separated near the edge of the piece of chromatography paper. This point of application is called the origin or baseline. The solvent front is the distance travelled by the solvent from the origin. Because of the absorbent nature of paper, the solvent moves against gravity and carries the ink dyes along with it. Stage 1: The ink dye is spotted and allowed to dry. The original spot is identified as A. The solvent begins to move up the paper by capillary action. Stage 2: Solvent moves up the paper taking different components along at different rates. Stage 3: The separation of the mixture is complete. The different components string out along the paper Note: * the water molecules attached to the chromatography paper form the stationary phase and the solvent moving through the paper is the mobile phase. * the stationary phase will tend to stop the components of the mixture from moving while the mobile phase will carry them along the paper. * the components of the mixture will be held to the stationary phase and will dissolve in the mobile phase to different extents * the components of the mixture will therefore move along the paper at different rates Each substance can be recognized by a value known as a retention factor (Rf) value, which is a constant for a given type of chromatography paper and a given solvent system. Rf = distance travelled by a component/distance travelled by solvent front 2.6) Extraction of Sucrose from Sugar Cane * Shredder: contains rotary knives for cutting cane into small pieces * Crusher: juice is extracted as water is sprayed on the cane and rollers apply pressure to crush the cane. The cane fibre, known as bagasse, is later burnt to supply fuel for the boilers * Clarifier: the cane juice is acidic and is neutralized by the addition of calcium hydroxide. Heating allows for the precipitation of impurities in the form of insoluble calcium salts * Filter: the precipitated solid materials are removed from the cane juice. The clarified juice passes to the first of several boilers * Boiler: the juice is evaporated under reduced pressure to prevent charring. Several boilers, each at a more reduced pressure than the previous one, are used * Crystallizer: the thick syrup is supersaturated. Pure sugar crystals are added to cause crystallization of the liquid * Centrifuge: here the mixture from the crystallizer is spun at high speeds to separate it into molasses and raw sugar crystals * Collectors: containers for collecting the separated molasses and sugar ________________ 3) Atomic Structure 3.1) An atom is the smallest part of an element that can exist and still show the properties of that element. Atoms contain subatomic particles called electrons, protons and neutrons. The protons and neutrons make up the nucleus of the atom. Electrons move around the nucleus of the atom in specific areas called shells. The shells are really different energy levels. Only electrons with the appropriate energy can occupy a given shell. The shell with the lowest energy level is closest to the nucleus and the energy level increases with each successive shell. Shells are represented as a series of concentric circles. The arrangement of electrons in an atom is called the electronic configuration. If the number of electrons in an atom is known, the electronic configuration can be worked out because the electrons in an atom fill the shells according to the following rules: * electrons fill up the shells in order, beginning with the lowest energy level * electrons generally fill one shell before entering the next shell The outer occupied shell of electrons in an atom is called the valence shell. The elections in the valence shell are the ones that take part in chemical bonding and are known as valence electrons. During chemical changes or reactions, these valence electrons are rearranged while the rest of the atom remains intact. 3.2) Relative mass of a particle is its mass compared with the mass of a proton. Relative charge on a particle is the charge compared with the charge on a proton. PROPERTIES OF ELECTRONS, PROTONS AND NEUTRONS Sub-atomic Particle Relative Mass Relative Charge Proton 1 +1 Neutron 1 0 Electron 0 -1 In each atom, the number of protons equals the number of electrons. A particle is neutral when the size of its positive charge equals the size of its negative charge. 3.3) The proton/atomic number of an atom is the number of protons in the nucleus. The nucleon/mass number of an atom is the sum of protons and neutrons present in the nucleus. An element is a substance made up of only one type of atom, i.e. atoms with the same atomic number. 3.4) Relative atomic mass is the average mass of one atom of an element compared to the mass of one atom of carbon-12 taken as exactly 12 units. ________________ a c 3.5) X where a - mass number, b - atomic number, c - charge, d - number of items in the entity b d X - symbol of element 3.6) Isotopes are atoms of the same element which have the same number of protons and different number of neutrons. They have identical chemical properties but may have slightly different physical properties. 3.7) Isotopes may occur naturally or be artificial. Artificial isotopes are made by bombarding the nuclei of atoms with neutrons or high-energy charged particles. Isotopes with unstable nuclei are radioactive. Radioactivity is the spontaneous disintegration of unstable atomic nuclei by the emission of radiation (alpha and beta particles, gamma radiation). Uses of Radioactive Isotopes * Diagnosing and Treating Disease * In medicine, radioisotopes are used to find out what is happening inside the body. Some substances concentrate naturally in some organs of the body. These are used as tracers, which are tiny quantities of radioisotopes used to study chemical changes within living organisms. e.g. Iodine-131 is used to measure the amount of radioactivity in the thyroid gland. A hyperactive thyroid will accumulate more iodine than a normal thyroid. Technetium-99 can be used to see if there is reduced blood flow - a possible sign of heart disease. * Radiation can also be used to treat diseases. Radiotherapy is used to treat cancer because it destroys cancer cells more readily than it destroys normal cells. A beam of radiation is directed at the cancerous tumour. e.g. Cobalt-60 produces gamma rays which are used to kill cancerous cells. Isotope Medical Uses Cobalt-60 treating certain types of cancer Sodium-24 tracing the flow of blood and to locate obstructions in the circulatory system Iodine-131 monitoring and treating goitre and other thyroid problems; also used in treatment of liver and brain tumours Thallium-201 monitoring certain heart diseases Technetium-99 monitoring certain heart diseases Plutonium-238 providing energy for heart pacemakers ________________ * Carbon-14 Dating * Carbon-14 dating is a method of finding out the age of archaeological specimens that were once living, such as bone, cloth, wood and plant fibres. It measures the levels of the isotope carbon-14 in samples of material that were once living. The instant an organism dies, the amount of carbon-14 decreases due to radioactive decay. * Energy Generation * Energy can be provided by unstable radioactive isotopes. Large radioactive atoms release energy when they are split. This is called nuclear fission. Fission of uranium-235 is a source of energy and electricity in some nuclear power stations, but must be monitored carefully. An accident could result in the exposure of humans and other living organisms to dangerous radiation., with disastrous consequences. The disposal of radioactive waste from nuclear power plants is also a problem that has not yet been adequately solved. * The human heart has a natural pacemaker that should maintain a regular heartbeat, but sometimes, it is defective and has to be replaced. Some artificial pacemakers are powered by the energy produced in radioactive decay. Plutonium-238 is used in thermoelectric batteries. As the plutonium decays, the heat produced is used to generate electricity that then stimulates the heart. 4) Periodic Table and Periodicity 4.1) The Periodic Table is an arrangement of the elements based on their atomic number into horizontal rows, called periods, and vertical columns, called groups.The rows are of such a length that elements with similar physical and chemical properties fall directly beneath one another. ________________ As we go across a period, proton numbers increase by 1. If we use the proton number to write the electronic configuration of an element, we can predict the period and group to which it belongs. * All the atoms in a particular period have the same number of shells occupied by electrons. * The number of electrons in the outer shell increases by 1 in consecutive elements across a particular period. * The number of shells containing electrons is the same as the period number. * All the atoms in a particular group have the same number of electrons in their outer shell. Atomic Radius Atomic radii decrease across a row in the periodic table due to an increase in the effective nuclear charge. Within each group, the atomic radius increases with the period number. The nucleus is positively charged and it attracts the negatively charged electrons. The greater the number of shells between the outer electrons and the nucleus, the less the outer electrons feel the effect of the nucleus. The bigger the charge of the nucleus, the greater its pull on the outer electrons, thus making the atomic radius smaller. As we go across a period, the nuclear charge increases but the number of electron shells remains the same. As a result, the attraction of the nucleus for the outer electrons increases and the atomic radius decreases. On the other hand, as we go down a group, there are more electron shells and the attraction between the nucleus and the outer electrons becomes weaker. The result is that the atomic radius increases. Ionization Energy/Electron Affinity Ionization energy is the energy necessary to remove an electron, which is endothermic. Electron affinity is the energy change when an electron is accepted by an atom, which can be endothermic or exothermic. First ionization energy increases across the period because of increasing nuclear charge and decreases down the group. The electrons in the outermost shell are more strongly bound to the nucleus due to increasing nuclear charge. The greater the negative value of the electron affinity, the greater the tendency of an atom to accept an electron. Electron affinity increases across a row and decreases down a column. Electronegativity The electronegativity is how strongly the nucleus of an atom attracts the electrons of other atoms in a bond. As the atomic radius decreases, the effect of the nucleus will be felt more strongly at the outside of the atom. Therefore, electronegativity increases across a period from left to right. As the atomic radius increases, the effect of the nucleus will be felt less at the outside of the atom. This means that electronegativity decreases down a group. Fluorine is the most electronegative element while francium is the least electronegative. Reactivity/Electropositivity increases down a group and decreases across a period. Periodicity is the recurrence of elements with similar properties at regular intervals in the Periodic Table. 4.2) Group I and II Elements All the elements in Groups I and II are metals, Group II are alkaline earth metals and Group I is alkali metals. Trends in Physical Properties down Group II 1. Atomic radius increases down Group II. As the number of shells containing electrons increases from magnesium to barium, the radius of the atom increases. Shielding of valence electrons by the electrons in filled inner shells also contributes to increasing atomic radius down the group. 2. Ionisation energy decreases down Group II. Metal atoms ionize by losing their valence electrons to form cations. Ease of ionisation indicates how easily an atom forms ions. The lower the ionization energy value, the more easily the metal ionises. 3. Group II elements have relatively high melting and boiling points. 4. There is a general increase in density from calcium to barium. The masses of the atoms increase down the group. 5. Electronegativity values are generally low and show a slight decrease down the group. As atomic radius increases down the group, the nucleus of the atom attracts electrons less readily. 6. Metallic Character increases down Group II. Trends in Chemical Properties of Group II Elements Magnesium Calcium Barium Reaction with oxygen or air reacts but only if heated reacts if heated in air reacts rapidly at room temperature Reaction with water a slow reaction; hydrogen is evolved; insoluble magnesium hydroxide is formed brisk reaction; hydrogen is evolved; sparingly soluble calcium hydroxide is formed very rapid reaction; hydrogen is evolved; soluble barium hydroxide is formed Reaction with dilute acids, e.g. HCl rapid reaction; hydrogen is evolved; a salt, e.g. MgCl2, is formed very rapid reaction; hydrogen is evolved; a salt is formed violent reaction; hydrogen is evolved; a salt is formed * Magnesium and calcium react with oxygen or air only if heated, whereas barium reacts rapidly at room temperature. * Magnesium only reacts very slowly with cold water, whereas calcium and barium react vigorously with cold water. Hot water or steam is needed for a more vigorous reaction with magnesium. * The reaction with dilute hydrochloric acid intensifies as the group is descended. 4.3) The Group VII Elements The halogens - fluorine, chlorine, bromine, iodine and astatine - are non-metallic elements found in Group VII of the Periodic Table. These non-metals are highly reactive elements, as can be predicted from their position to the far right of the Periodic Table. Physical Properties of the Halogens The atoms of the halogens each contain seven valence electrons. They then share the single (unpaired) valence electron to form diatomic covalent molecules in which the atoms are held by strong single covalent bonds. Weak forces then exist between the molecules. The halogens therefore: * have low melting and boiling points * are more soluble in non-polar solvents than in water * are non-electrolytes * exist as gases, volatile liquids or soft solids at room temperature, depending on the relative strengths of the intermolecular forces. ________________ PHYSICAL PROPERTIES OF HALOGENS Element Atomic Radius/nm Physical State at rtp and atp Colour Melting Point/°C Boiling Point/°C Electronegativity Solubility in water/ mol/dm3 Chlorine 0.099 gas yellow green -113 -35 3.0 1.5 Bromine 0.114 liquid deep red with deep red vapour -0.07 59 2.8 4.2 Iodine 0.133 solid black with a sheen; gives purple vapour on sublimation 114 184 2.5 2×10-2 Trends as the Group is descended Explanation atomic radius increases the number of shells occupied by electrons increases electronegativity decreases as the atoms get larger, the ability to attract electrons decreases melting and boiling points increase as the size of the halogen molecules increases, the strength of the intermolecular forces also increases physical state changes from gas to liquid to solid as intermolecular forces increase, molecules are pulled closer together Chemical Reactivity of the Elements Fluorine, the first element of the group, is the most reactive non-metal known and its chemistry can be complicated by the fact that it will oxidise many substances, including water. Unlike metals, the reactivity of non-metals depends on the ease with which they gain electrons, that is their electronegativity. The halogens are highly electronegative and are therefore very reactive. Smaller atoms gain electrons more readily than larger ones and therefore, the reactivity decreases as the group is descended. The Halogens as Oxidising Agents The halogens readily accept electrons from the substances with which they react. Substances that accept electrons are oxidising agents. The ability of the halogens to attract electrons decreases down the group. This means that their ease of ionisation decreases, so we can deduce that their oxidising strength also decreases down the group. When the halogens act as oxidising agents, the coloured halogen molecules are changed to colourless halide ions. These colour changes help us to recognize when the halogens are acting as oxidising agents. The relative strength of the halogens as oxidising agents is reflected in displacement reactions. (Experiment 6.1) X2 + 2KY —————> 2KX + Y2 halogen X salt of halogen Y salt of halogen X halogen Y (coloured) (colourless) (colourless) (coloured) ________________ 4.4) Period 3 Elements Changes in Physical Properties Sodium Magnesium Aluminium Silicon Phosphorus Sulphur Chlorine Argon Atomic Number 11 12 13 14 15 16 17 18 Electron Configuration 2,8,1 2,8,2 2,8,3 2,8,4 2,8,5 2,8,6 2,8,7 2,8,8 Brief Description A soft metal of low density. Sources are rock salt, salt mines and sea water. Obtained by electrolysis of molten common salt. Silvery grey metal. Important in plant and animal life. Silvery grey metal. The most abundant metal. Sources are bauxite and aluminosilic-ate rocks. A non-metal. Has the diamond structure. Non-metallic solid. Widely distributed in the Earth’s crust. An essential constituent of cell protoplasm, the nervous system and bone. Yellow solid. Occurs near volcanoes, hot springs, in natural gas and petroleum. A green-yellow gas. Prepared by electrolysis of concentrate-d brine. An inert gas extracted from the atmosphere, especially during production of liquid air or the manufacture of ammonia. Used where an inert atmosphere is required. Metal or non-metal? metal metal metal semi-metal non-metal non-metal non-metal non-metal Atomic radius/nm 190 145 118 111 98 88 79 71 Electronegati-vity 0.9 1.2 1.5 1.8 2.1 2.5 3.0 0 Melting point/°C 97 649 660 1410 44 123 -113 -189 Boiling point/ °C 882 1090 2467 2355 280 444 -35 -186 ________________ Trends in Chemical Properties Reactions with Oxygen Reactions with Chlorine Reactions with Water Sodium When heated, sodium forms an oxide: 4Na(s) + O2(g) —> 2Na2O(s) ionic compound; basic oxide Δ 2Na(s) + Cl2(g) —> 2NaCl(s) ionic compound; crystalline solid This is a vigorous to violent reaction; hydrogen is liberated; the resultant solution is alkaline:2Na(s) + 2H2O(l) —> 2NaOH(aq)+ H2(g) Magnesium Magnesium burns with a brilliant flame, producing magnesium oxide: Δ 2Mg(s) + O2(g) —> 2MgO(s) ionic compound; basic oxide Δ Mg(s) + Cl2(g) —> MgCl2(s) ionic compound; crystalline solid Reacts slowly with hot water to form the metallic oxide and hydrogen gas: Mg(s) + H2O(l) —> MgO(s) + H2(g) Aluminium A thin oxide coating forms: 4Al(s) + 3O2(g) —> 2Al2O3(s) shows both ionic and covalent character; amphoteric oxide Δ 2Al(s) + 3Cl2(g) —> 2AlCl3(s) shows both ionic and covalent character; crystalline solid Reacts with steam to form aluminum oxide and hydrogen gas: 2Al(s) + 3H2O(g) —> Al2O3(s)+’3H2(g) Silicon No reaction Powdered silicon combines with chlorine: Si(s) + 2Cl2(g) —> SiCl4(s) covalent compound; colourless liquid No reaction Phosphorus Phosphorus forms two oxides: P4O6 and P4O10 Δ P4(s) + 3O2(g) —> P4O6(s) covalent compound; acidic oxide Phosphorus ignites in chlorine, burning with a pale yellow flame: Δ P4(s) + 6Cl2(g) —> 4PCl3(l) covalent compound; colourless liquid No reaction Sulphur Sulphur forms two oxides: SO2 and SO3 Δ S(s) + O2(g) —> SO2(g) covalent compound; acidic oxide Chlorine passed over molten sulphur yields covalent chlorides slowly S8(s) + 4Cl2(g) —> 4S2Cl2(l) No reaction Chlorine Chlorine forms a variety of covalent oxides Not applicable Dissolves in and reacts with water to form a mixture of acids: Cl2(g) + H2O(l) —> HCl(aq) + HOCl (aq) ________________ 5) Structure and Bonding A compound is a substance formed when two or more different types of elements combine chemically. Atoms combine by forming bonds. A chemical bond is a force of attraction between combining atoms. During chemical bonding: * Atoms use their outer shell or valence electrons to form bonds. Valence electrons are the electrons in the outer electron shell of an atom. * Atoms may lose, gain or share their valence electrons. * As these changes occur, the electron configurations of the atoms change and new particles are formed. (These are ions or molecules). Ions are charged particles formed when atoms lose or gain electrons. Molecules are particles formed when atoms combine by sharing electrons. * Elements combine to acquire a stable electronic configuration, i.e. a filled valence shell. Metal atoms containing 1,2 or 3 valence electrons tend to lose electrons. The larger the atomic radius of the atom, the more easily it loses electrons. Some non-metal atoms with 5, 6 or 7 valence electrons may gain a number of electrons to fill their valence shell. The smaller the non-metal atom, the more readily it accepts electrons. Non-metal atoms containing 4 to 7 valence electrons may also share electrons when combining with other non-metals. There are three main types of chemical bonds: 1. the electrovalent or ionic bond 2. the covalent bond 3. the metallic bond 5.1, 5.2) Ionic Bonds Ionic bonds are the attractive forces that hold oppositely charged ions together in electrovalent compounds, i.e. when a metal reacts chemically with a non-metal. ‘Dot and cross’ diagrams are used to illustrate the formation of bonds. Only the valence electrons are shown for the combining atoms. Although all electrons are the same, dots (・) and crosses (×) are used to indicate which atom the electrons came from. In the bonding between sodium and chlorine: * Each sodium atom transfers its valence electron to the valence shell of a chlorine atom. * As a result of this transfer, the sodium atoms are changed to positive ions (cations). The charge on the cation is equivalent to the number of electrons lost. * Each chlorine is changed to a negative ion (anion). The charge on the negative ion is negative to the number of electrons gained. * The cations and anions formed have new electronic configurations, where they had full valence shells. * Attractions between the oppositely charged ions provide the binding forces which hold ionic compounds together. * The smallest part of an ionic compound is known as the formula unit. The formula unit of sodium chloride (NaCl) consists of one sodium ion and one chloride ion. Some more examples of ionic bonding are the formation of potassium fluoride (KF), magnesium chloride (McCl2) and lithium oxide (Li2O). Note: All valence electrons transferred from the metal must be accepted by the non-metal. This means that the ratio of metal to non-metal is variable. Covalent Bonds Covalent bonds are formed when atoms of non-metallic elements combine with one another. * The non-metal atoms share one or more pairs of electrons. * Each non-metal atom shares a number of electrons so that each atom appears to have the electronic configuration of the nearest noble gas in the Periodic Table. * Each shared pair of electrons will constitute a covalent bond. * The terms single, double and triple covalent bonds are used to describe the sharing of one, two or three pairs of electrons, respectively, between a pair of atoms. * When non-metal atoms share electrons, the particles formed are called molecules. Molecules are particles formed when two or more atoms combine by covalent bonds. The atoms in a molecule may be the same or different. * The attraction between the nuclei of the atoms and the shared pair(s) of electrons provides the binding force which holds the atoms together. Covalent Bonds between Non-Metal Atoms of the Same Element Fluorine (F2): Each isolated fluorine atom contains unpaired electrons. These isolated electrons form a pair of shared electrons in the fluorine molecule. The shared electron pair represents a single covalent bond. Each fluorine atom in a fluorine molecule contains three lone pairs. Lone pairs or non-bonded pairs are pairs of electrons found in molecules that are not involved in the formation of simple covalent bonds. Oxygen (O2): Each isolated oxygen atom contains two unpaired electrons. These isolated electrons form two pairs of shared electrons in an oxygen molecule. The two shared electron pairs represent two covalent bonds, known as a double bond. Each oxygen atom in an oxygen molecule contains two lone pairs. Nitrogen (N2): Each isolated nitrogen atom contains three unpaired electrons. These isolated electrons form three pairs of shared electrons in a nitrogen molecule. The three shared electron pairs represent three covalent bonds, known as a triple bond. Each nitrogen atom in a nitrogen molecule contains one lone pair. ________________ Covalent Bonds between Different Atoms Water (H2O): In water, an oxygen atom is singly bonded to two hydrogen atoms. There are two lone pairs of electrons on the oxygen atom. Ammonia (NH3): In ammonia, a nitrogen atom is singly bonded to three hydrogen atoms. There is one lone pair of electrons on the nitrogen atom. Methane (CH4): In methane, a carbon atom is singly bonded to four hydrogen atoms. There are no lone pairs of electrons on the central (carbon) atom. All the valence electrons of carbon are used in the bonding in the methane molecule. Carbon Dioxide (CO2): In carbon dioxide, a carbon atom is doubly bonded to two oxygen atoms. Two electron pairs are shared between each oxygen atom and the carbon atom, giving two double bonds within the molecule. Co-ordinate Covalent Bonding A co-ordinate or dative covalent bond is formed when both electrons in the shared pair come from the same atom. e.g. ammonium chloride, hydroxonium ion, carbon monoxide, nitric acid Non-polar and Polar Covalent Molecules When a covalent bond is formed between two different types of atoms, the atoms attract the electron pair to different extents. Atoms of fluorine, oxygen, nitrogen and chlorine attract electrons more strongly than atoms of other elements, and are therefore strongly electronegative. In molecules such as hydrogen chloride, the electron pair of the bond is pulled closer to the more electronegative chlorine atom. As a result of this, the chlorine atom develops a tiny negative charge, whereas the hydrogen atom develops a tiny positive charge. This molecule has a slight separation of charge within it and is described as a polar molecule. A polar molecule is a covalent molecule that has areas with different charges due to unequal sharing of electrons. By contrast, a molecule of hydrogen (H2) or chlorine (Cl2) is non-polar since the atoms are identical and have the same attraction for the bonding electron pair. Examples of highly polar molecules include hydrogen fluoride (HF), water (H2O) and ammonia (NH3). Note: All polar molecules are electrically neutral overall since the sum of the small positive and negative charges equals zero. Polar covalent compounds have higher melting and boiling points and react differently with water than non-polar molecules do. (Experiment 4.1) 5.3) The molecular formula of a compound gives the actual numbers of the different types of atoms in one molecule of a covalent compound or the ratio of ions present in one formula unit of an ionic compound. COMMON CATIONS Monovalent Divalent Trivalent Hydrogen (H+) Magnesium (Mg2+) Iron (III) (Fe3+) Lithium (Li+) Calcium (Ca2+) Aluminium (Al3+) Sodium (Na+) Barium (Ba2+) Chromium (III) (Cr3+) Potassium (K+) Iron (II) (Fe2+) Copper (I) (Cu+) Copper (II) (Cu2+) Silver (Ag+) Zinc (Zn2+) Ammonium (NH4+) Tin (II) (Sn2+) Lead (II) (Pb2+) COMMON ANIONS Monovalent Divalent Trivalent Fluoride (F-) Oxide (O2-) Nitride (N3-) Chloride (Cl-) Sulphide (S2-) Phosphate (PO43-) Bromide (Br-) Carbonate (CO32-) Phosphide (P3-) Iodide (I-) Sulphite (SO32-) Hydride (H-) Sulphate (SO42-) Hydroxide (OH-) Dichromate (VI) (Cr2O72-) Nitrite (NO2-) Peroxide (O22-) Manganate (VII) (MnO4-) Chromate (VI) (CrO42-) Hydrogensulphate (HSO4-) Ethanedioate (C2O42-) Hydrogencarbonate (HCO3-) Ethanoate (CH3COO-) Nitrate (NO3-) Methanoate (HCOO-) Cyanate (OCN-) Cyanide (CN-) Hypochlorite (ClO-) Chlorite (ClO2-) Chlorate (ClO3-) Perchlorate (ClO4-) Thiocyanate (SCN-) Naming Ionic Compounds A simple method to work out formulae is to swap valencies. Place the valency of the first element after the symbol of the second and place the valency of the second element after the symbol of the first. e.g. Magnesium Nitride Valencies of elements: Mg - 2 N - 3 Chemical Formula: Mg3N2 When naming compounds containing an element which can have more than one valency (transition metals), the valency is shown by Roman numerals in brackets after the name of an element. e.g. Copper (I) Oxide - Cu2O vs. Copper (II) Oxide - CuO Naming Covalent Compounds Prefix Number Indicated mono 1 di 2 tri 3 tetra 4 penta 5 hexa 6 hepta 7 octa 8 nona 9 deca 10 Rules: * The first element is named first, using element’s name. * Second element is named as an anion. * Prefixes are used to denote the number of atoms * “Mono” is not used to name the first element. e.g. SO2 - Sulphur Dioxide SO3 - Sulphur Trioxide N2O - Dinitrogen Monoxide NO - Nitrogen Monoxide N2O4 - Dinitrogen Tetroxide ________________ 5.4) Metallic Bonding Metal atoms do not bond together by ionic bonds or by covalent bonds. In a metal, the valence electrons leave the atoms, which then form positive ions. The electrons lost by a particular atom do not necessarily remain associated with the resulting ion. In fact, the electrons are mobile and they flow through the spaces between the positive ions. Metals can be viewed as orderly arrangements of positive ions held together in a ‘sea’ of freely moving electrons. This type of bonding, known as metallic bonding, is used to explain how the particles in metallic elements are held together. The metallic bond is therefore the strong electrostatic force of attraction which exists between the stationary positive ions and the mobile electrons in the metal. 5.5) Ionic Compounds Ionic solids are formed by ionic bonding - the particles in the solids are ions. These solids are crystals because of their structure. The regular structure of a crystal tells you that the ions are arranged in an orderly manner. We call this orderly arrangement a crystal lattice. A lattice is a regular arrangement of points in two or three dimensions. A crystal lattice is the arrangement of points on which atoms, molecules or ions are centre in a crystal. Ionic compounds are said to consist of ‘a giant structure of ions’. Simple Molecular Substances Most substances that contain covalent bonds are also described as simple molecular substances because they consist of separate molecules. Covalent substances tend to be liquids or gases at room temperature because the forces between their particles are weak. Intermolecular forces are weak forces between individual molecules (such as Van der Waals forces). Intramolecular forces are strong forces within molecules (covalent bonds) which hold the atoms together in the molecule. Intermolecular Forces * Van der Waals Forces All covalent molecules, whether polar or non-polar, develop temporary or instantaneous dipoles. This results from the uneven movement of all the electrons within the molecules. Van der Waals forces are the weak attraction between oppositely charged ends of molecules within temporary dipoles. * Hydrogen Bonds The hydrogen bond is the weak attraction between the slightly positive hydrogen atom in one polar molecule and the slightly electronegative atom in another polar molecule of the same type or of a different type. Giant Molecular Crystals Like simple covalent solids, these solids are also formed by covalent bonding, but they do not form individual molecules. Instead, they exist as macromolecules in which very strong covalent bonds extend in three dimensions. Giant covalent solids may contain one type of atom, as in the case of diamond and graphite, or more than one type of atom, as in the case of silicon dioxide. These solids have high melting and boiling points as well as high heats of fusion and vaporization Structure of Diamond Structure of Graphite ________________ 5.6) PROPERIES OF IONIC AND COVALENT COMPOUNDS Property Ionic Compounds Simple Covalent/Molecular Compounds Composition Ions held together by strong ionic bonds. Molecules with strong covalent bonds between atoms, but weak forces between molecules. State at Room Temperature Crystalline solids - owing to strong ionic bonds holding ions together in a 3-dimensional lattice. Most are liquids or gases due to weak forces between molecules. Melting and Boiling Points, Heats of Fusion and Vaporization High - strong ionic bonds require a lot of energy to separate. Low - weak intermolecular forces require little energy to break. Solubility Most are soluble in water but insoluble in covalent (organic) solvents. Most are soluble in organic solvents, e.g. ethanol, but insoluble in water Electrical Conductivity Conduct electricity when molten or dissolved in water - ions are free to move Do not conduct electricity in any state - no free ions or electrons are present 5.7) Sodium Chloride Crystals * white in colour; transparent * cubic owing to arrangement of ions * brittle, easily breakable * no lubricating property * fairly high melting point because of strong ionic bonds * conducts electricity when molten or dissolved in water because of free ions Diamond Graphite Each atom is bonded to four others. Atoms are arranged in flat six-membered rings. It is one of the hardest substances known. It is widely used in cutting and drilling. Soft and slippery because of weak binding. Colorless, reflects colors of other substances; shiny, long-lasting, can be cut with diamonds, sparkly. Light can be reflected through the diamond. Opaque, dark grey, has a sheen to it Due to the absence of free electrons, diamond does not conduct electricity. Conducts electricity because of the presence of free electrons in its structure. Insoluble in water and organic solvents Insoluble in water and organic compounds No lubricating property The lubricating properties of graphite can be explained in terms of the ability of the flat sheets to slide past each other. High melting point High melting point due to strong covalent bonds 5.8) Allotropes are different forms of the same element existing in the same physical state. e.g. diamond and graphite are allotropes of carbon. 6) Mole Concept 6.1) Relative atomic/formula/molecular mass is the average mass of one atom/formula unit/molecule of an element/compound compared to the mass of one atom of carbon-12 taken as exactly 12 units. A mole is the amount of substance which contains the number of particles as there are atoms in 12 grams of carbon-12, i.e. Avogadro’s number (6.02×1023) of particles. The molar mass of a compound or element is the relative atomic/molecular mass expressed in grams. 6.2) Calculations involving the Mole Number of moles is equal to given mass divided by the molar mass. Given mass = # of mols × molar mass # of mols = Given mass/molar mass # of particles = Avogadro’s Number × # of mols (Experiment 8.1) 6.3) Avogadro’s Law states that equal volumes of gases under the same conditions of temperature and pressure contain the same number of molecules. The molar volume of a gas is in the volume that contains one mole of molecules of the gas. * At standard temperature and pressure (s.t.p.), which is a temperature of 273K/0°C and a pressure of 101kPa/1 atm, the volume of 1 mole of any gas is 22.4dm3. * At room temperature and pressure (r.t.p.), which is a temperature of 298K/25°C and a pressure of 101kPa/1 atm, the approximate volume of 1 mole of any gas is 24dm3. 6.4) The law of constant composition states that all pure samples of the same chemical compound contains the same elements combined together in the same proportions by mass. A chemical formula shows how many moles of each element combined to form one mole of the compound. 6.5) Writing and Balancing Equations A balanced equation has the same number of atoms of each element on both sides of the equation. e.g. 3Mg(s) + N2(g) ——> Mg3N2(s) CuSO4(aq) + 2NaOH(aq) ——> Cu(OH)2(s) + Na2SO4(aq) Rules for Writing Equations * Use chemical symbols and formulae to describe what happens in the reaction. List all reactants on the left of the arrow, and all products on the right. * Represent elements by their symbols. For solid elements, use the symbol for a single atom, for example Mg, Fe, C, etc. For gases, use the symbol for the molecule, for example most common elements are diatomic - thus H2, O2, N2, etc. * Use state symbols (s), (l), (g) and (aq) to indicate the physical state of the substances in the equation. * Check symbols and formulae to ensure that they are all correct. * Balance the equation by inserting the whole numbers in front of each term of the equation where necessary. Do not change the numbers within the formula. Ionic Equations Ionic equations are equations that show only the species taking part in a reaction between ionic substances. How to Balance Ionic Equations * Complete and balance the equation in the form in which it is given. If the equation is given in the molecular form, then complete it in molecular form, balance it and then change it to an ionic equation. * Write all strong electrolytes in the ionic form. Write other species in the molecular form. * Ensure that there is the number of each type of atom or polyatomic ion on both sides of the equation. * Omit all ions appearing on both sides of the equation unchanged. * Write the ionic equation. Include only those ions that have undergone a change. * Ensure that the net charges are equal on both sides. 6.6) (Chapter 10) 6.7) Mass concentration is the mass of the solute (in grams) dissolved in 1000cm3 (1dm3) of solution. Molar concentration is the number of moles of solute dissolved in 1000cm3 (1dm3) of solution. Volumetric Analysis One of the common experimental methods for determining the concentrations of solutions is known as volumetric analysis. The main aspect of volumetric analysis is measuring the volume of one solution of accurately known concentration which is required to react quantitatively with a solution of the substance being determined. The solution of accurately known concentration is called the standard solution. The standard solution is usually added to a titration flask via a pipette. The solution whose concentration is to be determined is then added to the standard solution from a burette. The process of adding solution from burette to solution in titration flask until the reaction is complete is termed a titration. The point at which reaction is complete is known as the equivalence point or end-point, and the volume of solution added from the burette is the titre. The apparatus commonly used in volumetric analysis includes a balance, a pipette, a burette and a volumetric flask. (Experiment 9.1) ________________ The Rough Titration Place the solution and indicator (if needed) in the titration flask. Then add the other solution from the burette, shaking continuously until there is a slight excess, as indicated by the first permanent change in the indicator. Note that color changes are best seen against a white background. Presentation of Titration Results Burette Readings / cm3 Rough Titration Titration 1 Titration 2 final burette reading initial burette reading volume of acid used average volume used = Acid-base Titrations An acid-base titration can be generally represented by the equation: HX + BOH —> BX + H2O acid base salt water Acid-base reactions are neutralization reactions. They may be represented by an ionic equation: H+(aq) + OH-(aq) —> H2O The general procedure for the titration of a soluble base (an alkali) with an acid is as follows: 1. The aqueous solution of the alkali is placed in the titration flask. 2. A few drops of a suitable indicator are added. 3. The acid is run into the titration flask from a burette. In practice, the acid in the burette is slowly added to the alkali until there is a colour change in the contents of the conical flask. The color change is due to the reaction of the acid with the indicator and is the end-point of the reaction. (Experiment 9.2) 7) Acids, Bases and Salts 7.1) An acid is a substance which, in solution, produces hydrogen ions (H+) as the only positive ions. In fact, though, the hydrogen ions produced do not remain separate in solution. They become attached to the oxygen atoms in the polar water molecule forming ‘hydronium’ as follows: H2O + H+ ——> H3O+ When acids come into contact with water then, they do not merely dissolve - chemical changes take place. Considerable heat is evolved when concentrated sulphuric acid is dissolved in water. This is an indication that new bonds are being formed. When not dissolved in water (in their anhydrous state), acids may be: * solids, e.g. citric acid and ascorbic acid (vitamin C) * liquid, e.g. nitric acid and phosphoric acid * gases, e.g. HCL (dissolves in water to form hydrochloric acid), SO3 (dissolves in water to form sulphuric acid) and NO2 (which dissolves in water to form a mixture of nitric acid and nitrous acid) Preparation of Acids An acid anhydride is a non-metallic oxide that dissolves in water to form an acid. SO2(g) + H2O(l) ——> H2SO3(aq) sulphurous acid (sulphuric (IV) acid) SO3(g) + H2O(l) ——> H2SO4(aq) sulphuric acid (sulphuric (VI) acid) CO2(g) + H2O(l) ——> H2CO3(aq) carbonic acid N2O5(g) + H2O(l) ——> 2HNO3(aq) nitric (V) acid 2NO2(g) + H2O(l) ——> HNO2(aq) + HNO3(aq) nitrous acid (nitric (III) acid) nitric (V) acid A base is a proton acceptor. Many bases are oxides and hydroxides of metals. However, any substance that accepts protons is a base. Bases are also defined as substances that react with acids to form salts and water only: base + acid ——> salt + water MgO(s) + 2HCl(aq) ——> MgCl2(aq) + H2O(l) 2KOH(aq) + H2SO4——> K2SO4(aq) + 2H2O(l) An alkali is a base that is soluble in water. Commonly used alkalis are NaOH, Ca(OH)2, KOH and a solution of ammonia in water (NH3.H2O). Salts are formed when metal ions or the ammonium ion take the place of the hydrogen ion (or ions) of an acid: NaOH(aq) + HCl(aq) ——> NaCl(aq) + H2O(l) acid salt An acidic oxide is an oxide of non-metal that neutralizes bases to form salt and water only. A basic oxide is an oxide of metal which neutralizes acids to form salt and water only. An amphoteric oxide is an oxide of metal that neutralizes both acids and bases to form a salt and water. A neutral oxide is an oxide of non-metal that neither neutralizes acids or bases. 7.2) The pH scale is a number scale which indicates whether a solution is alkaline, acidic or neutral. Acidic solutions have a pH of less than 7, neutral solutions have a pH of 7 and alkaline/basic solutions have a pH of more than 7. The stronger the acid, the lower the pH. The stronger the base, the higher the pH. (Experiment 11.1) ________________ 7.3) Acids that dissociate/ionize completely are strong electrolytes and strong acids. Acids that dissociate partially are weak electrolytes. Common Strong Acids Anions of these Strong Acids Formula Name Formula Name HCl Hydrochloric acid Cl- Chloride ion HBr Hydrobromic acid Br- Bromide ion HI Hydroiodic acid I- Iodide ion HNO3 Nitric acid NO3- Nitrate ion HClO4 Perchloric acid ClO4- Perchlorate ion HClO3 Chloric acid ClO3- Chlorate ion H2SO4 Sulphuric acid HSO4- SO42- Hydrogen Sulphate ion Sulphate ion Common Weak Acids Anions of these Weak Acids Formula Name Formula Name HF Hydrofluoric acid F- Fluoride ion CH3COOH Acetic/Ethanoic acid CH3COO- Acetate/Ethanoate ion HCN Hydrocyanic acid CN- Cyanide ion HNO2 Nitrous acid NO2- Nitrite ion H2CO3 Carbonic acid HCO3- CO32- Hydrogen Carbonate ion Carbonate ion H2SO3 Sulphurous acid HSO3- SO32- Hydrogen Sulphite ion Sulphite ion H3PO4 Phosphoric acid H2PO4- HPO42- PO43- Dihydrogen Phosphate ion Hydrogen Phosphate ion Phosphate ion (COOH)2 Oxalic acid H(COO)2- (COO)22- Hydrogen Oxalate ion Oxalate ion Bases can be strong or weak depending on the extent to which they dissociate and produce OH- ions in solution. Most metal hydroxides are strong electrolytes and strong bases. Ammonia (NH3) is a weak electrolyte and weak base. 7.4) Characteristic Reactions of Acids * They give particular colours with indicators. With universal indicator, the exact colour depends on the pH of the acid * With active metals, acids react to give hydrogen gas and a salt: metal + acid ——> salt + hydrogen Zn(s) + H2SO4(aq) ——> ZnSO4(aq) + H2(g) * Acids react with metal oxides (and hydroxides) to form a salt and water only: CuO(s) + H2SO4(aq) ——> CuSO4(aq) + H2O(l) NaOH(aq) + HNO3(aq) ——> NaNO3(aq) + H2O(l) This is a neutralization reaction. * Acids react with carbonates and hsydrogencarbonates to yield a salt, water and carbon dioxide: Na2CO3(s or aq) + 2HCl(aq) ——> 2NaCl(aq) + H2O(l) + CO2(g) NaHCO3(aq) + HNO3(aq) ——> NaNO3(s) + H2O(l) + CO2(g) (Experiment 11.2) 7.5) Vitamin C (Experiment 27.4) Vitamin C (chemical name: ascorbic acid) has the molecular formula C6H8O6. It is obtained from a variety of sources, e.g. West Indian cherries, blackcurrants, citrus fruits, cabbage and spinach. Here are some of the roles of Vitamin C: * it controls the formation of dentine, cartilage and bone * it helps in the formation of red blood cells and the healing of wounds Vitamin C is water-soluble and is easily oxidised. The extent of oxidation is increased by: * cutting or crushing - this releases enzymes called oxidases * increasing the temperature * the action of alkalis * traces of copper Losses of vitamin C occur when fruits and vegetables are stored. Further losses occur during preparation and cooking because vitamin C breaks down on cooking. Sodium hydrogen carbonate (baking powder) is sometimes added to boiling vegetables as it helps them keep their green colour. Although this improves the appearance, it reduces the vitamin C content because the alkaline hydrogen carbonate reacts with ascorbic acid by neutralisation: HCO3-(aq) + H+(aq) ——> CO2(g) + H2O(l) Methanoic Acid Methanoic acid, also called formic acid, has the molecular formula HCOOH. It is the simplest of the carboxylic acids and is used in processing textiles and leather. It is an important intermediate in chemical synthesis and occurs naturally, most notably in some ants. Lactic Acid 2-Hydroxypropanoic acid, also called lactic acid, has the molecular formula C3H6O3. Formation of lactic acid n muscles occurs when insufficient oxygen is supplied to the muscles, resulting in the release of energy via anaerobic cellular respiration. Lactic acid is a toxic chemical that can cause muscles to stop working and its presence is sometimes described as an ‘oxygen debt’. Acetic Acid in Vinegar Acetic acid or Ethanoic acid has the molecular formula CH3COOH. It has a low pH level, therefore, it is highly acidic and destroys bacteria. It is used in food preservation. Lime Juice Lime juice contains citric acid, which reacts with iron oxide (rust) so that it can be removed. 7.6) Reaction of Bases with Ammonium Salts Ammonium salts are decomposed when mixed with a strong base to produce ammonia gas. NaOH(aq) + NH4Cl(aq) ——> NaCl(aq) + H2O(l) + NH3(g) NaOH(aq) + NH4NO3(aq) ——> NaNO3(aq) + H2O(l) + NH3(g) The ammonia is readily detected by its pungent odour and by turning damp red litmus blue. The ionic equation for the reactions is: NH4+(aq) + OH-(aq) ——> H2O(l) + NH3(g) 7.7) Salts Solubility Rules Nitrates All nitrates are soluble Chlorides, Bromides and Iodides (Halides) All halides are soluble, except the halides of silver and lead; lead chloride and lead bromide, however, are soluble in hot water Sulphates All sulphates are soluble, except barium sulphate and lead sulphate; calcium sulphate and silver sulphate are slightly soluble Carbonates All carbonates are insoluble, except sodium carbonate, potassium carbonate and ammonium carbonate Note: All common salts containing the ions Na+, K+ and NH4+ are soluble. ________________ Salt Colour and other Characteristics Uses Ammonium Chloride / NH4Cl white crystals dry cells (batteries), fertilisers Ammonium Sulphate (sulphate of ammonia) / (NH4)2SO4 white crystals fertilisers Calcium Carbonate (marble, limestone) / CaCO3 white but can be coloured decorative stones, manufacture of cement and lime Calcium Sulphate (plaster of Paris, gypsum) / CaSO4 white crystals plastering walls, making casts, etc. Magnesium Sulphate (Epsom salts) / MgSO4 white crystals purgative Copper (II) Sulphate / CuSO4 blue crystals fungicides Sodium Carbonate (washing soda) / Na2CO3 white crystals or powder in cleaning, in laundry as a water softener, in the manufacture of glass Ionic Precipitation Ionic precipitation reactions are sometimes referred to as double displacement decomposition reactions. In ionic precipitation reactions, two soluble ionic compounds react to form one soluble and one insoluble compound. Ionic precipitation reactions can be represented generally by the equation: AB + CD ——> AD + CB The procedure is as follows: * the aqueous solutions are mixed * the mixture is then warmed, if necessary, and filtered * the residue is washed and dried Examples of ionic precipitation reactions are: AgNO3(aq) + NaCl(aq) ——> AgCl(s) + NaNO3(aq) BaCl2(aq) + H2SO4(aq) ——> BaSO4(s) + 2HCl(aq) FeSO4(aq) + 2NaOH(aq) ——> Fe(OH)2(s) + Na2SO4(aq) Test for Completeness of Precipitation * Add precipitating agent to the filtrate. * If cloudy, the reaction is incomplete. If not cloudy, the reaction is complete. Direct Combination In direct combination reactions, two or more elements react to form a single product. Direct combination reactions can be represented by the general equation: A + B ——> AB where A and B are elements. The following are examples of direct combination reactions. * The reaction of a metal with oxygen to produce the oxide of the metal: metal + oxygen —> metal oxide 2Mg(s) + O2(g) —> 2MgO(s) Some metals combine with the oxygen of the air at room temperature; others have to be heated to form the oxides. * Heating a non-metal in air or oxygen to produce the non-metal oxide: non-metal + oxygen —> non-metal oxide Δ S(s) + O2(g) —> SO2(g) * The reaction of a metal with a non-metal to produce a salt: metal + non-metal —> salt 2Na(s) + Cl2(g) —> 2NaCl(s) Fe(s) + S(s) —> FeS(s) Substitution (Displacement) Reactions In substitution or displacement reactions, one element displaces another element from a compound. Displacement reactions can be represented by the general equation: A + BX —> AX + B A and B may be metals or non-metals. The action of acids on excess metal, insoluble metal oxide or metal hydroxide, metal carbonate or hydrogencarbonate: Zn(s) + H2SO4(aq) —> ZnSO4(aq) + H2(g) PbO(s) + 2HNO3(aq) —> Pb(NO3)2(aq) + H2O(l) Mg(OH)2(s) + 2HNO3(aq) —> Mg(NO3)2(aq) + 2H2O(l) CaCO3(s) + 2HCl(aq) —> CaCl2(aq) + H2O(l) + CO2(g) 2NaHCO3(s) + H2SO4(aq) —> Na2SO4(aq) + 2H2O(l) + 2CO2(g) It is common practice whenever a solid and an acid are used in such preparations to use excess of the solid. This ensures that all the acid is used up in the reaction. Under these conditions, the acid is the limiting reagent. On completion of the reaction, excess (unreacted) solid is removed by filtration. A sample of the salt crystals can then be obtained from the filtrate by carrying out the following steps: * concentrate the filtrate by gentle evaporation (use a water bath) * cool the concentrate, testing for crystal formation - slow cooling leads to the formation of big crystals * filter to collect the crystals * carefully wash and dry the crystals 7.8) Cement Cements are materials used in buildings and other constructions to bind aggregates such as sand and crushed stone into a solid mass. Cements are used to produce concrete and mortars. Concrete is probably used on a larger scale than any other artificial material. When concrete is freshly mixed, it is plastic and malleable so that it can be cast in almost any shape. In construction, it can be used both as a structural element, since it is strong and durable, and as a decorative element, since it can be coloured, painted and decorated. How is cement made? The most important type of cement produced today is Portland cement. The approximate chemical composition of Portland cement is 60-70% calcium oxide, 17-25% silica, 3-8% aluminum oxide and small amounts of other materials. The raw materials used to make cement include limestone, clay, and sand. (Experiment 31.1) ________________ Chemical Used to Preserve Mechanism of Action Comments Nitrites and Nitrates: NaNO2 and NaNO3 heat-processed meat, poultry and fish, cheeses control the growth and toxin production of Clostridium botulinum Nitrites produce nitric oxide, which gives a bright red colour to meats; Nitrites combine with other products to form nitrosamines, which have been found to be a health hazard to humans (cancer causing) Sulphur Dioxide and Sulphites: SO2 and Na2SO3 soft fruits, fruit juices, lemon juices, beverages, wines, pickles and fresh shrimps antioxidants - they prevent browning in fresh and dried fruits and vegetables; sulphates also act as antimicrobials Some people can be allergic; products have to be appropriately labelled Ascorbic acid (Vitamin C / C6H8O6) , Citric acid / C6H8O7 bread, cakes, cheeses, pickles antioxidants; inhibit enzymes by lowering the pH Gives an acidic taste to food; can replace Vitamin C lost in processing Acetic acid (in vinegar) (CH3COOH) bread, cakes, cheeses, pickles anti-microbial, inhibit bacterial and fungal growth; low pH provides a hostile medium for microbes Loss of colour and change in flavour of the food; not suitable for people with certain ailments Sodium Chloride (common salt) used, for example, in curing meats and fish anti-microbial, restricts the growth of bacteria; reduces the amount of available water and alters osmotic pressure Salt hardens the flesh of fish and meat and decreases its ability to absorb moisture; high sodium intake is linked to high blood pressure Wood smoke (by burning woods such as hickory, mahogany, etc); liquid smoke processed meat and fish one reason is for flavour; smoke also has chemicals with antibacterial properties, e.g. phenols, formaldehyde Smoke also contains carcinogenic chemicals; it is suggested that the risk of colon cancer can be reduced by minimizing the intake of barbecued or smoked foods Spices and condiments, e.g. cinnamon and cloves meats and fish mainly used for flavouring; anti-microbial in small concentrations Onion and garlic seem to have potential use as anti-microbial agents Sugar (C12H22O11) fruits, candies high osmotic pressure and low water levels prevents microbial growth Too much sugar may damage teeth; not suitable for diabetics 7.9) Acid salts are formed when only some of the H+ ions are replaced. Acid salts, therefore, contain some H+ ions from the original acid. Only dibasic and tribasic acids can form acid salts. NaOH(aq) + H2SO4(aq) —> NaHSO4(aq) + H2O(l) Normal salts are formed when all of the H+ ions are replaced. 2NaOH(aq) + H2SO4(aq) —> Na2SO4(aq) + 2H2O(l) 7.10) All neutralisation reactions involve reaction between an acid (or acidic oxide) and a base. Water is produced in neutralisation reactions; heat is also given out, i.e. neutralisation reactions are exothermic. Neutralisation reactions can be investigated in the laboratory by following the change in pH or in temperature as the acid is added to the alkali. (Experiment 12.1) Indicators are dyes that are one colour in acidic solution and another colour in a solution of a base. Litmus Methyl Orange Screened Methyl Orange Phenolphthalein Colour with acid red red light red colourless Colour with water purple orange grey colourless Colour with base blue yellow green pink ________________ Teeth Structure of Teeth * enamel - this is the hardest of all animal tissues and it covers the crown of the tooth * dentine - forms the bulk of the tooth; it is harder than bone but not as hard as enamel * pulp - soft tissue which contains the blood vessels, nerves and the cells which form dentine Why take care of your teeth? It is essential that we keep our teeth clean to avoid a buildup of plaque and to avoid tooth decay. Plaque is a sticky, filmy layer of bacteria that constantly forms on the teeth, below and above the gums. The bacteria in plaque can produce toxins that affect gums, causing the gums to become red and swollen. If not attended to, this can lead to loss of many teeth. If plaque is not removed, it may become hard and is called calculus/tartar. Acids are produced in the mouth when food particles are broken down. Bacteria living in plaque also produces acids. These acids attack calcium compounds in the teeth and the teeth decay. The compounds are calcium carbonate (CaCO3) and calcium hydroxyphosphate (Ca5(OH)(PO4)3), more commonly known as apatite. CaCO3(s) + 2H+(aq) —> Ca2+(aq) + CO2(g) + H2O(l) Toothpaste Brushing with toothpaste removes plaque and food particles which tend to collect around the teeth, thereby reducing the incidence of tooth decay. Brushing also helps to remove stains and leaves the mouth fresh. Toothpastes are designed to scrub away the plaque, harden the teeth enamel and have a pleasant taste. Toothpastes must also contain materials that allow the ingredients to stay together as a gel or paste that stays on the toothbrush. Ingredient Function Example Abrasive To dislodge the plaque. Leaves the teeth feeling smooth. Should be mild so as not to damage the teeth. Sodium bicarbonate; calcium carbonate Fluoride To harden the teeth enamel. It is believed that fluoride ions replace hydroxyl ions in apatite: Ca5(OH)(PO4)3 —> Ca5F(PO4)3 hydroxyapatite fluoroapatite Fluoroapatite is more resistant to attack by acids in the mouth since it does not contain OH- ions. Sodium monofluorophosphate Detergents Make the toothpaste foam and helps to remove fatty films Sodium laurel sulphate Humectants Minimise water loss from the toothpaste. Aid the retention of flavours Glycerol; sorbitol Thickeners or binding agents Prevent the ingredients from separating out Cellulose Flavours These are carefully blended. Menthol, mint, spearmint, clove Preservatives Prevent mould and bacteria from spoiling the toothpaste Sodium benzoate 7.11) Volumetric Analysis Calculations ________________ 8) Oxidation-Reduction Reactions 8.1) Bleaching Agents Bleaching agents are compounds which are used to remove color from substances such as textiles. In earlier times textiles were bleached by exposure to the sun and air. Today most commercial bleaches are oxidizing agents, such as sodium hypochlorite (NaOCl) or hydrogen peroxide (H2O2) which are quite effective in "decolorizing" substances via oxidation. The decolorizing action of bleaches is due in part to their ability to remove these electrons which are activated by visible light to produce the various colors. The hypochlorite ion (OCl-), found in many commercial preparations, is reduced to chloride ions and hydroxide ions forming a basic solution as it accepts electrons from the colored material as shown below. OCl- + 2e- + HOH --------> Cl- + 2 OH- Browning of Cut Fruit Apples and other produce (e.g., pears, bananas, peaches, potatoes) contain an enzyme (called polyphenol oxidase or tyrosinase) that reacts with oxygen and iron-containing phenols that are also found in the apple. The oxidation reaction basically forms a sort of rust on the surface of the fruit. You see the browning when the fruit is cut or bruised because these actions damage the cells in the fruit, allowing oxygen in the air to react with the enzyme and other chemicals. The reaction can be slowed or prevented by inactivating the enzyme with heat (cooking), reducing the pH on the surface of the fruit (by adding lemon juice or another acid), reducing the amount of available oxygen (by putting cut fruit under water or vacuum packing it), or by adding certain preservative chemicals (like sulfur dioxide). On the other hand, using cutlery that has some corrosion (as is seen with lower quality steel knives) can increase the rate and amount of the browning by making more iron salts available for the reaction. Rusting Rusting is an oxidation process. Using iron, it loses electrons according to the equation: Fe(s) —> Fe2+(aq) + 2e- This reactions occurs in parts of the iron that are not exposed to air. In parts of the iron where there is a good supply of oxygen and water, the following reaction occurs: O2(g) + 2H2O(l) + 4e- —> 4OH-(aq) The Fe2+ ions react with the OH- ions to form iron (II) hydroxide: Fe2+ + 2OH- —> Fe(OH)2(s) Oxygen from the air oxidises the iron (II) hydroxide to hydrated iron (III) oxide, which is rust: Fe(OH)2(s) + O2(g) —> Fe2O3..𝑥H2O(s) (Experiment 13.2) 8.2) 8.3) Rules for Determining Oxidation States/Numbers * In all uncombined elements, atom’s oxidation number equals 0. * In all compounds, sum of oxidation numbers equals 0. * In all ions, sum of oxidation numbers equals ion charge. * In all compounds: Group 1 elements - +1 Group 2 elements - +2 Group 3 elements - +3 Fluorine - -1 Oxygen - -2 Chlorine - -1 * In a binary covalent compound, the more electronegative atom is given the negative oxidation number and the less electronegative atom is given the positive oxidation number. * In most compounds H=+1, except when bonded to a metal, e.g. HCl, NaH * In most compounds, O=-2, except when bonded to F or in peroxides (O=-1). 8.4) Identifying Oxidation and Reduction Reactions Half equations are equations that show the separate oxidation and reduction processes in any redox reaction. e.g. 2Ca(s) + O2(g) —> 2CaO oxidation: Ca(s) —> Ca2+(aq) + 2e- reduction: O2(g) + 4e- —> 2O2-(aq) 8.5) Oxidising and Reducing Agents An oxidising agent brings about the oxidation of another substance. In the process, it is reduced. A reducing agent brings about the reduction of another substance. In the process, it is oxidised. Oxidising Agent Colour Change or other observable sign Products Concentrated Nitric Acid (HNO3) brown gas evolved NO2(g), water, nitrate(aq) Hot Concentrated Sulphuric Acid (H2SO4) gas produced; has characteristic smell SO3(g), sulphate(aq), formed with metals or their compounds Potassium Manganate (VII) (KMnO4) / Dilute Sulphuric Acid (H2SO4) colour changes from purple to colourless Mn2=(aq) Potassium Dichromate (VI) (K2Cr2O7) / Acid colour changes from orange to green Cr3+(aq) Iron (III) salts: Fe3+(aq) yellow to pale green Fe2+(aq) Hydrogen Peroxide (H2O2) effervescence, colourless oxygen gas evolved O2 I2(aq) from brown to colourless I-(aq) Reducing Agent Colour Change or other observable sign Products Hydrogen Sulphide (H2S) yellow colloidal suspension sulphur and water Sulphur Dioxide (SO2) no significant observable change H2SO4 or SO42-(aq) Sulphite (SO32-) no observable change H2SO4 or SO42-(aq) Concentrated HCl yellow-green gas chlorine (Cl2) KI / H+(aq) brown solution or black precipitate of I2 iodine, water Fe2+(aq) turns yellow or brown Fe3+(aq) Hydrogen Peroxide (H2O2) effervescence, colorless hydrogen gas evolved H2(g) ________________ 8.6) Reactant Observable Change Equation KI(aq) solution changes from colourless to brown (iodine) 2I-(aq) —> I2(aq) + 2e- Fe2+(aq) colour change from pale green to yellow brown Fe2+(aq) —> Fe3+(aq) + e- Na2S or H2S colloidal suspension of sulphur S2-(aq) —> S + 2e- Tests for Reducing Agents A reducing agent should decolourise acidified potassium manganate (VII) ion: MnO4-(aq) + 5e- —> Mn2+(aq) purple from reducing agent colourless A reducing agent should also change acidified potassium dichromate (VI) (orange) to the green chromium (III) ion: Cr2O72-(aq) + 3e- ——> 2Cr3+(aq) orange from reducing agent green 9) Electrochemistry 9.1) Electrical Conductors Electrical conductors allow a current to flow through them. Insulators do not allow the passage of an electric current. (Experiment 13.1) Electrical conductors can be divided into two groups: metals and graphite and electrolytes. Electrolytes are compounds that conduct an electric current and are decomposed by it. Metallic conductors (and graphite) Electrolytes Conduct in the solid and liquid states Do not conduct in the solid state; Conduct in the molten state or in aqueous solutions No chemical changes occur when they conduct electricity Chemical changes occur when they conduct an electric current 9.2) Metallic Conduction Conductivity in metals (and graphite) is due to the presence of mobile electrons. Mobile electrons are present in the solid metal and graphite. The metals remain chemically unchanged by the passage of the current. Electrolytic Conduction Conductivity in electrolytes is due to the presence of mobile ions. The positive and negative ions are separated, and it is this separation that results in the decomposition of the electrolyte. Ionic compounds do not conduct electricity in the solid state, because the ions are rigidly and tightly held in a crystal lattice. In the solid state, the ions are not free to move. In the molten state or in solution, the lattice breaks down and the ions are free to move. Electrolytes are molten or aqueous ionic compounds and also some polar covalent compounds which react with water to produce ions. 9.3) Strong and Weak Electrolytes Electrolytes may be classified as strong or weak, based on how completely they ionize. * Strong electrolytes are completely ionized in aqueous solution and so, have a high concentration of ions in the electrolyte. Examples are strong acids (H2SO4), strong bases (NaOH), aqueous or molten ionic compounds (NaCl). * Weak electrolytes are partially ionized in aqueous solutions so have a relatively low concentration of ions in the electrolyte. Examples are water, weak acids (CH3COOH), weak bases (aqueous NH3). 9.4) Electrolysis is the decomposition of an electrolyte by the passage of an electric current through it. Electrodes are the points where current enters and leaves an electrolyte. They are collecting rods, often made of graphite or platinum, that are inert and usually do not take part in the chemical changes occurring during electrolysis. Some electrodes are active and these do take part in these chemical changes. For example, platinum reacts with chlorine gas.and graphite reacts with oxygen gas. The positive electrode, or anode, is connected to the positive terminal of the power supply or battery. The negative electrode, or cathode, is connected to the negative terminal of the power supply or battery. During electrolysis, anions (negative ions) move towards the anode and cations (positive ions) move towards the cathode. 9.5) Molten Aqueous BaCl2 Ba2+ 2Cl- NaOH Na+ OH- AgBr Ag+ Br- NaCl Na+ Cl- Al2O3 2Al3+ 3O2- KCl K+ Cl- PbBr2 Pb2+ 2Br- H2SO4 2H+ SO42- Aqueous solutions contain additional H+ and OH- ions from water, totaling 4 ions in the solution: 2 from the electrolyte and 2 from water. Only one cation and one anion are discharged. The electrolysis of aqueous solutions uses the theory of selective discharge. 9.6, 9.7, 9.8) Discharge of Ions During electrolysis, there is discharge of ions. Discharge is the process by which ions gain or lose electrons and become atoms or molecules. Ions lose their charge during discharge. Cathode Anode Movement of Ions Positive ions (cations) move towards the cathode Negative ions (anions) move towards the anode Electrons Cations gain electrons from the cathode Anions lose electrons to the anode Result Cations become neutral atoms or molecules Anions become neutral atoms or molecules Half-Equation Mn+ + ne- —> M Xn- —> X + ne- Reduction or Oxidation? Reduction - gain of electrons Oxidation - loss of electrons Preferential Discharge Ease of Discharge Cation Product at Cathode Anion Product at Anode Difficult Easy K+ Hydrogen from Water SO42- Oxygen from water Na+ NO3- Ca2+ Cl- Chlorine Mg2+ Br- Bromine Al3+ I- Iodine Ni2+ Nickel OH- Oxygen Pb2+ Lead H+ Hydrogen Cu2+ Copper Ag+ Silver At the cathode: * In concentrated solutions of nickel or lead compounds, nickel or lead will be discharged, instead of hydrogen ions of water, which is less reactive than nickel or lead. * In very dilute solutions, hydrogen, copper and silver ions are preferable to be discharged, according to the ease to be discharged. * Reactive ions (K+, Na+, Ca2+, Mg2+, Al3+, etc) will never be discharged in either concentrated or dilute conditions. Instead, hydrogen ions from water will be discharged at the cathode. At the anode: * In concentrated solutions, iodide, chloride and bromide ions are preferable to be discharged, although it is harder to discharge compared with hydroxide ions. * In very dilute solutions containing iodide, chloride and bromide ions, hydroxide ions of water will be discharged, instead of the halides mentioned, according to the ease of discharge. * Sulphate and nitrate are never discharged in concentrated or dilute conditions. Electrolysis using Different Types of Electrodes Inert electrodes, such as platinum and graphite, do not react with electrolytes or products during electrolysis. Active electrodes, such as sodium and copper, react with electrolytes and products of electrolysis, affecting the course of electrolysis. 9.9, 9.10) Electrolysis Calculations The mass of a substance produced at the electrodes (or consumed at reactive anode) during electrolysis is proportional to: * the electric current (in amperes) * the time (in seconds) over which a constant current passes ________________ The electric charge in coulombs (C) transferred in electrolysis is given by: Q = It charge = electric current × time Faraday Constant, F This is the quantity of electric charge carried by one mole of electrons or one mole of singly charged ions. Its approximate value is 96500 C/mol. When we electrolyze silver nitrate (AgNO3) solution using silver electrodes, silver is deposited at the cathode. Ag+(aq) + e- —> Ag(s) One Faraday (96500 C) is required to deposit one mole of silver. This is the same amount of electricity that is required to remove one mole of silver from the silver anode. When we electrolyze copper sulphate (CuSO4) solution using copper electrodes, copper is deposited at the cathode. Cu2+(aq) + 2e- —> Cu(s) In this case, it requires two Faradays to deposit one mole of copper. This is because two moles of electrons are needed to produce one mole of copper atoms from one mole of copper (II) ions. 9.11) Applications of Electrolysis in Industry Electrolysis is used commercially in a number of ways. Not only does it play a part in extraction of some elements, but it is very useful in enhancing the appearance of metals and protecting them from corrosion. Corrosion occurs when a metal reacts with substances in the environment forming oxides and sometimes sulphide, carbonates, hydroxides and sulphates. Examples of the use of electrolysis include: * extraction of reactive metals such as sodium and aluminum from their ores or compounds * extraction of active non-metals such as the halogens * electroplating, e.g. chrome plating, nickel plating and galvanizing * anodizing aluminium * electrorefining, e.g. in obtaining pure copper from impure copper Anodizing Aluminium Anodizing is the process by which aluminum is given a thick protective coat of aluminum oxide by electrolysis. When exposed to the air, aluminum combines with oxygen to form aluminum oxide (Al2O3). The oxide forms an even coat and seals the surface, thus protecting the metal from further corrosion. Electrolysis is used to make this protective layer thicker and tougher. Aluminum is made the anode of an electrolytic cell that contains dilute sulphuric acid or dilute chromic (VI) acid as the electrolyte. (Any electrolyte that releases oxygen gas at the anode can be used). The liberated oxygen combines with the aluminum anode coating it with oxide. The reaction at the anode is: 4OH-(aq) —> 2H2O(l) + O2(g) + 4e- Another advantage of anodizing is that this protective layer can be made to absorb dyes, which then are permanently fixed by treatment with boiling water. Electroplating Electroplating is the process of covering one metal with a protective coat of another metal by electrolysis. In practice, the outer layer is made of a less reactive metal.There are several reasons for electroplating metal objects, including: * to enhance the appeal of the plated article - the electroplated article is more decorative than the metal underneath * to protect the covered metal from corrosion * to avoid using expensive metals for the object In electroplating, the object to be covered is made the cathode and the pure plating metal is made the anode. The solution needs to contain ions of the plating metal. During plating, ions pass into solution from the pure anode and are discharged as a thin layer on the cathode. Electrorefining Electrorefining is a process in which the purity of metals, such as copper, is improved by electrolysis. Copper of high quality is needed for some purposes, such as wiring. This is because electrical resistance increases with the quantity of impurities in the copper. Copper obtained by methods such as chemical reduction is not sufficiently pure. * Electrolyte: mixture of copper (II) sulphate and sulphuric acid * Anode: impure copper * Cathode: a strip of pure copper Anode half-equation: Cu(s) —> Cu2+(aq) + 2e- Copper atoms leave the anode and enter the solution as copper ions. Cathode half-equation: Cu2+(aq) + 2e- —> Cu(s) Copper ions are discharged at and deposited on the cathode. Overall, copper leaves the anode and is deposited on the cathode. The main impurities include zinc and gold. Zinc atoms in the impure copper anode behave like copper atoms, i.e. they lose electrons and pass into solution as zinc ions. The gold and other unreactive impurities collect at the bottom of the anode as a solid mixture, which is referred to as anode mud. 10) Rates of Reaction 10.1) Rate of reaction is the change in concentration of reactant or product with time at a stated temperature Conditions Required for Chemical Reactions to Occur Chemical reactions involve the breaking of bonds (in the reactants) and the formation of new bonds (in the products). Chemists believe that for this to happen, three conditions are necessary: * Reactant molecules must collide. * Reactant molecules, on collision, must have energy equal to or greater than the necessary activation energy. Activation energy is a certain minimum energy required for bonds within reactant molecules to break and for the particles to become sufficiently energized for products to be formed. * Reactant molecules must collide in the correct position (i.e. with the correct orientation). Collisions which are sufficiently energetic would be most effective if the colliding molecules approached each other in such a way that the energy released on collision can be passed on directly to the bonds to be broken. Collisions of particles in any other position will not result in the formation of products. ________________ Effective collisions are those collisions that result in the formation of products. They require that all the reactants are correctly oriented and have the required activation energy. 10.2, 10.3, 10.4) Factors which Affect the Rates of Chemical Reactions Concentration The theory is that an increase in concentration means there are more reactant molecules in a given volume. This increases the chances of more frequent effective collisions and thus leads to a faster reaction rate. (Experiment 14.1) Pressure Only gaseous reactions are affected as gas is compressible. At higher pressure, molecules are forced to move closely together, hence increasing the particles per unit volume of gas and effectively increases the collisions between reacting molecules so the speed of reaction increases. High pressure is commonly used in industrial processes so that the reaction goes faster. Temperature Speed of reaction increases when temperature increases. Particles do not always react upon collision but just bounce as they do not have enough activation energy to react. With an increase in temperature, particles absorb the energy and, having enough activation energy, they move faster and collide more effectively and frequently per second. Therefore, the speed of reaction is increased. We store food at the lower temperature of the refrigerator to reduce the rate of spoilage reactions. We cook food in pressure cookers because water boils at about 120°C in pressure cookers and the food cooks faster. (Experiment 14.2) Catalysts Catalysts are chemical substances which alter speeds of reaction without being used up at the end of a reaction. They can be reused and only a small amount of a catalyst is needed to affect a reaction. Catalysts lower the need for energy to break bonds so activation energy is lower. Consequently, bond breaking occurs more easily and more often when particles collide. A catalyst in the solid state may catalyze reactions in the gaseous state or in aqueous solution by providing a surface for the reactants to react. Enzymes are biological catalysts. They are very specific; one enzyme catalyzes one type of reaction. They are sensitive to temperature; they work best at 40°C. Too high or too low temperatures destroy enzymes. They are also sensitive to pH; they function within a narrow range of pH. (Experiment 14.3) Surface Area The rate of a chemical reaction involving a solid reactant is increased by increasing the state of subdivision of the solid, while decreasing the surface area of the solid has the opposite effect. In such reactions, collisions occur between moving molecules and the solid reactants. It follows that the smaller the particles of the solid, the greater the surface area available for collisions, and the greater the reaction rate. 11) Energetics 11.1) Exothermic and Endothermic Reactions (Experiment 15.1) Exothermic reactions are chemical changes that result in an increase in the temperature of the surroundings Examples of exothermic changes include: * the neutralization reaction between aqueous sodium hydroxide and hydrochloric acid * dissolving solid sodium hydroxide in water * burning propane gas * respiration - a reaction occurring in living systems in which energy stored in the chemical bonds of food is released Endothermic reactions are chemical changes that result in a fall in the temperature of the surroundings. Endothermic reactions are much rarer than exothermic reactions. Examples of endothermic changes include: * the dissolving of some substances in water, e.g. ammonium nitrate, potassium iodide, urea and sodium thiosulphate * photosynthesis * the reaction between steam and carbon The kinetic energy and the chemical energy together make up the energy if a chemical (also called its heat content). This energy is represented by the symbol H. Energy changes occur during the course of chemical reactions, as reactants form new products. The energy change during a reaction is represented by ΔH. When the energy change (ΔH) occurs at constant pressure, it is also described as the enthalpy change. ΔH = Hproducts - Hreactants If Hproducts > Hreactants : * overall heat is absorbed in the reaction * the reaction vessel becomes cold * ΔH is positive * bonds are broken * the reaction is endothermic If Hproducts < Hreactants : * overall heat is released in the reaction * the reaction vessel becomes hot * ΔH is negative * bonds are formed * the reaction is exothermic 11.2) Energy Profile Diagrams The minimum energy which must be supplied before reactions proceed is known as the activation energy of the reaction. In all chemical reactions, old bonds must be broken before new ones can be formed. For this reason, reactants must be supplied with energy. Catalysts work by providing a pathway with lower activation energy. 11.3) Measuring Enthalpy Changes When carrying out experiments to measure heat changes in the laboratory, the following items of apparatus are required: * an insulated container to serve as a calorimeter (the apparatus to measure heat evolved or absorbed in a chemical reaction), e.g. styrofoam cup * a thermometer * a balance * volumetric apparatus such as pipette and/or burette or measuring cylinder Some general steps in the procedure are: * allowing a known mass or volume of reactants to reach the steady temperature of the surroundings - this temperature is then recorded * thoroughly mixing the reactants and recording the highest or lowest temperature reached * determining the temperature change (Δθ) for the reaction * calculating the heat evolved or absorbed in the experiment * calculating the enthalpy change for the reaction enthalpy change = mass × specific heat capacity × temperature change ΔH = mcΔT ΔT = initial temp - final temp Specific heat capacity is the quantity of heat in Joules required to raise the temperature of unit mass or volume of a substance by 1°C/K. The specific heat capacity of water is 4.2J/g/K or 4200J/kg/K. Enthalpy of Neutralization (Experiment 15.2) ΔHn is the energy change which occurs when 1 mole of water is formed by the reaction of an acid and a base. e.g. 50cm3 of 2mol/dm3 HCl was placed in a styrofoam cup. 50cm3 of 2mol/dm3 NaOH was added to the acid. The initial temp was 25°C and the final temp was 38.7°C. Calculate the heat of neutralization of the reaction between NaOH and HCl. NaOH(aq) + HCl(aq) —> NaCl(aq) + H2O(l) Assume 1cm3= 1g, m= 50 + 50 ΔH= mcΔT = 100g = 100 × 4.2 × -13.7 ΔT = 25°C - 38.7°C = -5754 J = -13.7°C 1dm3 ——> 1000cm3 2mol ——> 1dm3 𝑥dm3 ——> 50cm3 𝑥mol ——> 0.05dm3 1000𝑥 = 50 𝑥 = 0.1mol of HCl and NaOH ∴ 0.1mol of H2O formed. 𝑥 = 0.05dm3 ΔHn = -5754J/0.1mol = -57540 J/mol = -57.54 kJ/mol