CHEMICAL EXTRACTION. 8]VeiZg. Introduction. In this chapter

Transcription

1 8]VeiZg ) CHEMICAL EXTRACTION Introduction Chemists use different forms of energy to extract elements from their compounds. In some cases, the chemical bonds that bind the elements are not very strong so only a small amount of energy is needed to separate the elements. In other cases, where the bonding is strong, considerable amounts of energy are required. In this chapter 4.1 Physical and chemical change page XXX Energy and chemical change page XXX 79 Figure 4.1 Iron Figure is extracted 4.1 from iron ore in a furnace. Molten iron Iron can is extracted be cast or from converted iron ore to in steel. a furnace. Considerable Molten energy iron can is be required cast or for converted these chemical to steel. processes. Considerable energy is required for these chemical processes. 74 THE CHEMICAL EARTH

2 )#& Remember Before beginning this section, you should be able to: relate change of state to the motion of particles as energy is removed or added identify when a chemical reaction is taking place by observing changes in temperature, the appearance of a new substance or the disappearance of the original substance. Key content By the end of this section, you should be able to: identify the differences between physical and chemical change in terms of rearrangement of particles summarise the differences between the boiling and electrolysis of water as an example of the difference between physical and chemical change gather information using firsthand or secondary sources to observe the electrolysis of water, analyse the information provided as evidence that water is a compound, and identify an application of the use of this reaction analyse and present information to model the boiling of water and the electrolysis of water, tracing the movements of and changes in arrangements of molecules gather and present information from first-hand or secondary sources to write equations to represent chemical reactions. 0(93)#!,!.$ #(%-)#!, #(!.'% Physical changes Chemists define a physical change as one that does not lead to the formation of new chemical substances. In comparative terms, physical changes involve small energy changes. Where physical changes involve a change of state, they can easily be reversed. Ice, for example, can be readily converted to liquid water by heating. Cooling will reverse the process and turn the water back to ice. In this case, no new substances have formed. Physical changes include: filtration: Muddy water can be easily separated into its two components by filtration (see figure 4.3). Mixing the separated components can readily reform the mixture. evaporation and distillation: Salt can be recovered readily from salt water by distillation. The evaporation and condensation processes are physical changes. Mixing the separated salt and water can readily reform the salt water. cutting, hammering and rolling: Metals can be cut into pieces or rolled to produce thin sheets. The processes of melting and solidification can rejoin the pieces. change of state: Melting and solidification, as well as evaporation and condensation discussed above, are physical changes. Melted butter quickly turns back to solid butter when placed in the refrigerator. The sublimation of dry ice into gaseous carbon dioxide is also a physical change. Figure 4.4 shows particle diagrams illustrating the concept of a physical change. In some physical changes, chemical bonds are broken or formed to achieve the change. For example, when copper is rolled, the metallic bonds are broken and reformed. The formation of salt crystals when salt water evaporates requires the formation of ionic bonds. In other cases, there are only weak physical forces between molecules. For example, when water solidifies, the water particles are held in the crystal lattice by physical forces of attraction. These intermolecular forces will be investigated in chapters sublimation: the process of a solid turning directly into a vapour without the formation of the liquid state. For example, iodine crystals sublime to form purple iodine vapour on heating. intermolecular forces: weak attractive forces between all types of matter. These forces are weaker than chemical bonds. Figure 4.2 The melting of ice to form water is an example of a physical change. It involves a change of state that can easily be reversed. CHAPTER 4 CHEMICAL EXTRACTION 75

3 (a) (b),arge PARTICLES MIXED WITH SMALL PARTICLES &ILTRATION 3MALL PARTICLES PASS THROUGH THE PORES OF THE FILTER PAPER Figure 4.3 Filtration is a physical process. (a) The mud in muddy water is separated from the water. (b) The larger particles of mud do not pass through the pores of the filter, but the smaller particles do. 3OLID PARTICLES ONLY VIBRATE 3OLID -ELTING,IQUID 0ARTICLES MOVE ABOUT (EAT Figure 4.4 Particle diagram showing physical change Chemical changes Chemical changes form new chemical substances. This process requires breaking and reforming chemical bonds. Significantly more energy is required for chemical changes than for physical changes. Also, reversal of chemical change may be quite difficult. A burning match is an example of a chemical change that cannot be reversed. Figure 4.5 illustrates some chemical changes using particle diagrams. Figure 4.5 A burning match is an example of a chemical change that cannot be reversed. 76 THE CHEMICAL EARTH

4 Reaction of lead oxide with carbon Lead oxide + Heat + Carbon Lead Reaction of hydrogen and iodine Hydrogen Figure 4.6 Particle diagrams of chemical change Carbon monoxide + Heat + Iodine Hydrogen iodide iron (III) oxide + carbon monoxide Some examples of chemical changes include: extracting a metal from a mineral: Iron metal can be extracted from iron (III) oxide in a blast furnace. The iron mineral is mixed with carbon (coke) and heated to high temperatures in a blast of air. The carbon undergoes a chemical change in air to form carbon monoxide, and heat energy is released. The carbon monoxide combines with the iron (III) oxide to produce molten iron and carbon dioxide. The overall chemical change is summarised by the following word and balanced chemical equations. iron + carbon dioxide coke: the black solid formed following distillation of coal. Coke is composed mainly of carbon. Fe 2 O 3 (s) + 3CO(g) 2Fe(l) + 3CO 2 (g) 4.1 PRACTICAL ACTIVITIES Investigating the electrolysis of water 4.4 DATA ANALYSIS Boiling and electrolysing water Figure 4.7 dissolving a metal in an acid: Magnesium metal is an active metal that reacts rapidly with dilute hydrochloric acid to produce hydrogen gas and a solution containing magnesium chloride. Heat is released in this reaction. The overall chemical change is summarised by the following word and balanced chemical equations. magnesium + hydrochloric acid magnesium chloride solution + hydrogen gas Mg(s) + 2HCl(aq) MgCl 2 (aq) + H 2 (g) electrolysis: the decomposition of a chemical substance (in solution or the molten state) by the application of electrical energy electrolysis of water: Water is not a good conductor of electricity but it can be made conductive by adding a little dilute sulfuric acid. If the acidified water is then electrolysed, it breaks down to form hydrogen and oxygen gases. The electrolysis causes a chemical change. Mixing the hydrogen and oxygen gases together will not produce water unless the mixture is ignited with a spark or flame. CHAPTER 4 CHEMICAL EXTRACTION 77

5 SYLLABUS FOCUS 7. WRITING BALANCED CHEMICAL EQUATIONS The following information summarises the steps involved in writing a balanced chemical equation. 1. Reactants are the chemicals that are allowed to react. Write them on the left-hand side of the arrow. 2. Products are the chemicals produced in the reaction. Write them on the right-hand side of the arrow. 3. Write the word equation for the reaction. 4. Use the valency rules to write the chemical formula under each reactant and product. 5. Check each side of the equation for atom conservation. 6. If the atoms are unbalanced, place coefficients in front of each formula so that they are balanced. Re-check that the atoms are now balanced. 7. Use standard abbreviations to write the physical state next to each reactant and product: (s) = solid; (l) = liquid; (g) = gas; (aq) = aqueous or dissolved in water. Example: Combustion of methane 1. Reactants: methane gas, oxygen gas 2. Products: carbon dioxide gas, liquid water 3. Word equation: methane + oxygen carbon dioxide + water 4. Write the correct formula for each substance. (Remember: CH 4 is a compound containing one carbon atom and four hydrogen atoms.) CH 4 + O 2 CO 2 + H 2 O 5. Check for atom conservation (number of atoms of each element). Reactants: C = 1; H = 4; O = 2 Products: C = 1; H = 2; O = 3 Therefore, atoms numbers do not balance. 6. Insert coefficients to balance the atoms; for example, the 2 in front of H 2 O represents two molecules of water. CH 4 + 2O 2 7. Re-check atom balance. CO 2 + 2H 2 O Reactants: C = 1; H = 4; O = 4 Products: C = 1; H = 4; O = 4 Atom balance has been achieved. The equation is balanced. 8. Insert physical states. CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(l) 4.1 QUESTIONS 1. Classify the following processes as physical or chemical changes. (a) Melting lead (b) Pulling copper to form long wires (c) Caramelising sugar by heating (d) Combustion of candle wax (e) Centrifuging whole blood 2. Jessica poured 50 ml of water into a beaker and heated it with a Bunsen burner flame. She noticed bubbles of a colourless gas rising through the water as it heated. On reaching boiling point, the water formed a colourless vapour. When a cold clock glass was held above the beaker, the water vapour condensed into a clear, colourless liquid. This liquid was collected and its physical properties tested and compared with the original water sample. They were found to be identical. Jessica took a further sample of liquid water and electrolysed it at platinum electrodes. A colourless gas evolved at the positive electrode. This gas was observed to make a flame burn more brightly. At the negative electrode, a colourless gas was formed that burnt explosively in air when ignited. (a) State the conclusions that can be made about water from the heating evidence alone. (b) Jessica wondered whether her observations could be used to conclude that water was an element. Would she be justified in such a conclusion? 78 THE CHEMICAL EARTH

6 3. Iron is grey in colour. Sulfur is yellow. Both are chemical elements. Bert placed powdered samples of iron and sulfur in a mortar and ground them together for several minutes until the colour was uniform. A magnet was then held above the mortar and Bert observed that the iron particles clung to the magnet leaving the yellow powder at the bottom of the mortar. (a) Classify the magnetic separation as a physical or chemical change. Explain your classification. (b) Explain whether or not a new chemical substance was formed after the two powders were ground together. 4. Write balanced chemical equations for each of the following word equations. (a) iron (II) oxide (s) + carbon monoxide (g) iron (molten) + carbon dioxide (g) (b) potassium (solid) + oxygen gas potassium oxide (solid) )#' Remember Before beginning this section, you should be able to: qualitatively describe reactants and products in decomposition reactions identify that a new compound is formed by rearranging atoms rather than creating matter. (c) silver carbonate (solid) silver + carbon dioxide gas + oxygen gas (d) aluminium + oxygen gas aluminium oxide (solid) (e) sodium (solid) + chlorine gas sodium chloride (solid) 5. Figure 4.8 shows a particle diagram of the interaction between sulfur and oxygen gas. Explain whether the process is a chemical change or physical change. Oxygen Sulfur Heat Figure 4.8 Particle diagram of an interaction between sulfur and oxygen %.%2'9!.$ #(%-)#!, #(!.'% Decomposition and synthesis Various forms of energy can be used to decompose chemical compounds or to synthesise compounds from elements of other compounds. In this section, we will examine some common examples of decomposition and synthesis reactions. Decomposition reactions decomposition: the process of breaking a compound down into its component elements or simpler compounds Various forms of energy, including heat, light and electricity, can be used to decompose compounds. Here are some everyday examples of decomposition reactions: Heat energy is used in our industrialised society to decompose minerals to produce metals in smelters. synthesis: the formation of a compound from its elements or a more complex compound from simpler compounds In nature, ultraviolet light energy decomposes ozone molecules into oxygen gas and oxygen radicals. This process is important in preventing most high-energy UV rays reaching the Earth s surface. Lightning initiates decomposition reactions in the atmosphere by providing electrical energy to various gas molecules. Airbags in cars contain the chemical sodium azide, which decomposes by detonation to produce a large volume of nitrogen gas to inflate the airbag during a crash. Here are some examples of decomposition reactions that can be performed in a laboratory. CHAPTER 4 CHEMICAL EXTRACTION 79

7 Key content By the end of this section, you should be able to: identify light, heat and electricity as the common forms of energy that may be released or absorbed during the decomposition or synthesis of substances, and identify examples of these changes occurring in everyday life explain that the amount of energy needed to separate atoms in a compound is an indication of the strength of the attraction, or bond, between them plan and safely perform a firsthand investigation to show the decomposition of a carbonate by heat, using appropriate tests to identify carbon dioxide and the oxide as the products of the reaction gather information using firsthand or secondary sources to observe the effect of light on silver salts and identify an application of the use of this reaction. 4.2 PRACTICAL ACTIVITIES Thermal decomposition of magnesium carbonate Examples: 1. Thermal decomposition of gold oxide When a sample of brown gold (III) oxide is heated over a Bunsen burner flame in a test tube, it readily decomposes to produce a sample of lustrous gold. The oxygen is evolved as oxygen gas. gold (III) oxide 2Au 2 O 3 (s) gold + oxygen gas 4Au(s) + 3O 2 (g) 2. Light decomposition (photolysis) of silver bromide Silver bromide is decomposed by light in the ultraviolet part of the spectrum. The white crystals darken as black grains of silver metal form. Bromine vapour is released in the process. silver bromide 2AgBr(s) silver + bromine vapour 2Ag(s) + Br 2 (g) 3. Electrolytic decomposition of molten lead (II) bromide Lead (II) bromide crystals melt at a relatively low temperature (373 C) to form a clear, colourless liquid. The liquid is heated to 400 C and electrolysed using inert electrodes; air is removed from the apparatus during electrolysis to prevent any other reactions. A brown vapour of bromine is evolved at the positive electrode and silvery globules of lead form at the surface of the negative electrode and sink to the bottom of the vessel. lead (II) bromide (liquid) PbBr 2 (l) DC power supply lead (liquid) + bromine vapour Pb(l) + Br 2 (g) 4.5 DATA ANALYSIS Investigating the purity of limestone by thermal decomposition Molten lead (II) bromide Air remov via pump Inert electrodes Electric heating mantle Figure 4.9 Electrolysis of lead (II) bromide Figure 4.10 Decomposition of sodium azide produces nitrogen gas, which inflates a car s airbag. Synthesis reactions Many chemical compounds can be readily synthesised from their elements in a laboratory. Heat energy is the normal form of energy used to initiate synthesis reactions, although some are initiated by light and electrical energy. Here are some everyday examples of synthesis reactions: u Lightning can supply sufficient energy to nitrogen and oxygen molecules in the atmosphere to initiate synthesis of nitric oxide. A similar process occurs in a car s engine, where the spark causes nitrogen and oxygen molecules to combine. 80 THE CHEMICAL EARTH

8 u Rust is an oxide of iron that forms when iron structures are exposed to oxygen in the air. u The ammonia industry synthesises ammonia directly by combining nitrogen and hydrogen gases at high temperatures and pressures over a catalyst. Here are some more examples of synthesis reactions. 4.3 PRACTICAL ACTIVITIES The effect of light on silver halides Examples: 1. Synthesis of iron (III) chloride using heat energy Iron wool can be heated briefly in a Bunsen flame and then placed in a gas jar of chlorine gas. The jar rapidly becomes filled with brown iron (III) chloride smoke. Moisture in the jar causes the iron (III) chloride to dissolve rapidly to form deep brown droplets. iron + chlorine gas 2Fe(s) + 3Cl 2 (g) iron (III) chloride 2FeCl 3 (s) 2. Synthesis of hydrogen chloride using light Hydrogen gas combines explosively with chlorine gas when the reaction mixture is exposed to light. The light provides the necessary energy to break the chemical bonds of the chlorine molecule. (Note: This reaction should not be performed in the school laboratory.) hydrogen gas + chlorine gas H 2 (g) + Cl 2 (g) hydrogen chloride gas 2HCl(g) Bond energies The amount of heat necessary to separate atoms in a compound is a measure of the strength of the attractive or bonding forces between oppositely charged ions in ionic compounds, or between atoms in covalent molecules. Ionic bonds To understand the different strengths of ionic bonds, we can compare the energy required to decompose some ionic oxides to produce 1 kg of the metal. Table 4.1 shows that more energy is required to extract magnesium from its oxide than zinc or silver. We can conclude that the ionic bonds between magnesium ions and oxide ions are much stronger than those between oxide ions and either silver or zinc ions. Table 4.1 Energy required to decompose ionic compounds Metal produced Compound decomposed Energy required to produce 1 kg of metal (kj) silver silver oxide, Ag 2 O 144 zinc zinc oxide, ZnO magnesium magnesium oxide, MgO Covalent bonds Table 4.2 compares the average bond energies of covalently bonded atoms. It shows that multiple covalent bonds are stronger than single CHAPTER 4 CHEMICAL EXTRACTION 81

9 covalent bonds. Thus, molecules such as nitrogen and carbon dioxide are quite thermally stable. In the atmosphere, nitrogen gas is essentially inert. Humans breathe in air but do not use the nitrogen so breathe it out again. Nitrogen-fixing bacteria are some of the few organisms on Earth that can use nitrogen directly. The high energy of a lightning flash can also break the bond between nitrogen atoms. Table 4.2 Average bond energies of some covalent bonds Covalent bond Bond energy ( J/bond) Compound in which bond is located O H 7.7 water H H 7.2 hydrogen gas O==O 8.3 oxygen gas C==O 13.3 carbon dioxide NxxN 15.7 nitrogen gas 4.2 QUESTIONS 1. When a photographic film is exposed to light it darkens. This darkening happens because A light decomposes the silver nitrate in the emulsion. B silver atoms are formed as silver ions accept electrons ejected from bromide ions. C bromine atoms lose electrons to form bromide ions. D light photons cause an electron to be transferred from a silver ion to a bromide ion. 2. White magnesium carbonate thermally decomposes to form white magnesium oxide and carbon dioxide gas. A student performs this experiment in the school laboratory. Select the statement that is true about this experiment. A There is a colour change as the magnesium carbonate is heated. B The pop test can be used to identify carbon dioxide as the gaseous product of the reaction. C Complete decomposition has occurred if the final white solid product does not effervesce when treated with drops of hydrochloric acid. D Magnesium oxide will decompose to magnesium and oxygen if the material is heated for too long. 3. A common synthesis reaction that occurs in the atmosphere during lightning storms is the reaction between nitrogen and oxygen. Select the balanced equation that correctly identifies the reaction that occurs. A N 2 (g) + O 2 (g) 2NO(g) B N 2 (g) + O 2 (g) NO 2 (g) C N(g) + O(g) NO(g) D N(g) + 2O(g) NO 2 (g) 4. A sample of blue-green copper (II) carbonate is heated strongly. It turns black and a colourless gas is evolved. The gas is passed into a beaker of clear limewater. The limewater goes milky white in colour. (a) Use the result of the limewater test to identify the gas evolved. (b) Predict the chemical formula of the black product of the decomposition reaction by writing a balanced chemical equation for the thermal decomposition reaction. (c) A g sample of the blue-green powder was heated so that it decomposed completely. The mass of the remaining black powder was 6.44 g. Calculate the percentage loss in weight during decomposition of the copper (II) carbonate. 5. Luigi heated a sample of sodium in a Bunsen flame until it melted and started to burn. He placed the burning sodium in a gas jar of chlorine. The sodium burnt with a bright yellow flame and clouds of a white, crystalline smoke were formed. 82 THE CHEMICAL EARTH

10 (a) State whether energy was released or absorbed in this reaction. (b) Identify the white, crystalline substance that was synthesised in this reaction. (c) Luigi collected the white substance and heated it in a crucible until it melted. He inserted electrodes into it and passed an electric current through the melted material. A silvery substance appeared at the negative electrode and a gas was released at the positive electrode. (i) State whether energy was released or absorbed in this electrolysis. (ii) Identify the new substances produced at each electrode. 6. (a) Zinc sulfide and magnesium sulfide are decomposed to form their elements, using kj of energy in each reaction. Use table 4.3 to compare the masses of zinc and magnesium that are formed, assuming that there is no energy wastage. (b) Identify the more stable sulfide. Justify your response. 7. Energy is supplied to break the covalent O H bonds in gaseous water molecules. Use table 4.2 to answer the following questions. (a) Identify the number of O H bonds in each water molecule. (b) Assuming no energy wastage, calculate the total amount of energy required. (c) Identify the products of this decomposition reaction. (d) Explain why these products readily recombine to form water. Table 4.3 Metal produced Compound decomposed Energy required to produce 1 kg of metal (kj) zinc zinc sulfide, ZnS magnesium magnesium sulfide, MgS SUMMARY Compounds are pure substances with fixed chemical compositions. Physical changes do not lead to the formation of new substances. A chemical change leads to the formation of new substances. Compounds can be separated into their component elements by chemical separation techniques. Energy is required for this process. Heat, light and electricity are common forms of energy used to decompose compounds into their component elements or to produce simpler compounds. Heat (and sometimes light) is released during the synthesis of compounds from their elements. The strength of bonds between atoms or ions can be estimated from the energy needed to decompose a compound into its component elements. CHAPTER 4 CHEMICAL EXTRACTION 83

11 PRACTICAL ACTIVITIES 4.1 PRACTICAL ACTIVITIES INVESTIGATING THE ELECTROLYSIS OF WATER Aim To use a voltameter to electrolyse acidified water Safety issues Wear safety glasses throughout this experiment. Ensure that water and chemicals do not come into contact with electrical wiring. Identify other safety issues relevant to this experiment by reading the method. Materials hydrogen voltameter with platinum electrodes DC power pack leads and alligator clips retort stand and clamps test-tube rack 2 medium test tubes 2 rubber stoppers wooden splint wax taper matches 200 ml water acidified with 50 ml of 2 mol/l sulfuric acid Method 1. Ensure that your safety glasses are on. 2. Assemble the apparatus as shown in figure Note which electrodes are connected to the positive and negative terminals of the DC power pack. 3. With the taps open, fill the arms of the voltameter with acidified water. Close the taps. Add some more acidified water to the reservoir. 4. Connect the leads to the voltameter using alligator clips. 5. Turn on the current and adjust the voltage to 4 6 V. Reservoir Taps Acidified water + DC power supply Figure 4.11 Electrolysis of water in a voltameter Gases collect Platinum electrodes 6. Observe the electrolysis and record your observations in your workbook. 7. Adjust to a higher voltage and observe the rate of electrolysis. 8. Continue the electrolysis until one arm is filled with gas. Note whether this is the tube containing the positive or negative electrode. Turn off the current. 9. Collect the gases in separate labelled test tubes. Place a test tube over the end of one arm of the voltameter and slowly open the tap to expel the gas. Stopper the tube and place it in the test-tube rack. Repeat with the other arm. 10. Invert the test tube containing gas from the positive electrode and test it with a glowing wooden splint. Record your observations. A positive test indicates oxygen. 11. Invert the test tube containing gas from the positive electrode and test it with a lighted wax taper. Record your observations. A positive test indicates hydrogen. 12. Clean up and return all equipment. 84 THE CHEMICAL EARTH

12 PRACTICAL ACTIVITIES Results and analysis 1. Construct a suitable table and tabulate your observations. 2. Compare the volumes of gases collected at the two electrodes. 3. Identify the gases in each tube on the basis of the tests performed. 4. Describe how the rate of electrolysis changed as the voltage increased. 5. Summarise the evidence that electrolysis is a chemical change and that water is a compound. 6. Gather and process second-hand data to identify a practical use for the electrolysis of water. Conclusion Write a brief conclusion for this experiment. 4.2 PRACTICAL ACTIVITIES THERMAL DECOMPOSITION OF MAGNESIUM CARBONATE Aim To investigate the thermal decomposition of magnesium carbonate Safety issues Wear safety glasses throughout this experiment. Take care not to touch the hot tube. Identify other safety issues relevant to this experiment by reading the method. Materials 5 large test tubes rubber stopper with glass delivery tube (see figure 4.12) bosshead, clamp and stand Bunsen burner 2 Petri dishes spatula powdered magnesium carbonate limewater 1 mol/l hydrochloric acid phenolphthalein indicator Method Ensure that your safety glasses are on. Part A 1. Place a scoop of magnesium carbonate in one test tube and half-fill another test tube with limewater. Set up the apparatus as shown in figure Heat the magnesium carbonate gently at first and then strongly using a blue Bunsen flame. Allow the gas evolved to bubble into the tube of limewater and observe any changes. After several minutes, and while still heating the magnesium carbonate, remove the tube of CHAPTER 4 CHEMICAL EXTRACTION 85

13 PRACTICAL ACTIVITIES Magnesium carbonate powder Retort stand Bunsen burner limewater and place it in a test-tube rack. This stops limewater entering the hot tube, due to pressure differences when heating ceases, and causing the tube to break. 3. Continue to heat the magnesium carbonate for about another 10 minutes. Turn off the Bunsen burner and allow the tube and contents to cool. Carbon dioxide evolved Figure 4.12 Thermal decomposition apparatus Limewater turns white as carbon dioxide reacts. Part B 1. When the heated tube from part A is cold, put a sample of the heated white powder in a Petri dish. 2. Put a small amount of unheated magnesium carbonate in another Petri dish. 3. Add drops of 1 mol/l hydrochloric acid to each dish. Record your observations and note any differences in the rate of fizzing. 4. Put a small amount of unheated magnesium carbonate in a clean test tube and a small amount of the heated solid in another. Add about 2 ml of water to each tube and mix. Add 4 drops of phenolphthalein indicator to each tube and note the differences in the colour of the indicator. Record all your observations. Results and analysis 1. Record all your observations in a suitable format. 2. The following word equation describes the thermal decomposition reaction: magnesium carbonate magnesium oxide + carbon dioxide (a) Describe the experimental evidence that carbon dioxide was formed on heating the magnesium carbonate. (b) Magnesium carbonate is a weak base in water. Magnesium oxide is a stronger base than magnesium carbonate. Phenolphthalein is pale pink in weakly basic solutions and deeper pink or crimson in solutions that are more basic. Describe the experimental evidence that magnesium oxide was produced in the reaction. (c) Use your observations from the experiment with hydrochloric acid to explain whether or not your sample of magnesium carbonate completely decomposed on heating. Justify your response. (d) Write a symbolic equation for the decomposition reaction. 3. When hydrochloric acid reacts with magnesium carbonate, there is a rapid effervescence as carbon dioxide is evolved. The word equation is: magnesium carbonate + hydrochloric acid magnesium chloride + water + carbon dioxide Write a balanced symbol equation for this reaction. Conclusion Write a brief conclusion for this experiment. 86 THE CHEMICAL EARTH

14 PRACTICAL ACTIVITIES 4.3 PRACTICAL ACTIVITIES THE EFFECT OF LIGHT ON SILVER HALIDE SALTS Aim To investigate the effect of sunlight and UV light on silver salts and to identify an application of this reaction Background In this experiment, you will make silver halide salts by precipitation from silver nitrate solution. The silver halides are silver chloride, silver bromide and silver iodide. Figure 4.13 Some silver salts decompose under UV light. These three samples of silver chloride were placed in Petri dishes. One sample (top) is being irradiated with UV light. The samples at front are the control (left) and one that has previously been exposed to UV (right). The latter has turned grey due to the formation of metallic silver. Safety issues Wear safety glasses throughout this experiment. Do not look at the bulb in the UV lamp. Do not allow silver salts to contact your skin. Identify other safety issues relevant to this experiment by reading the method. Materials 3 small Petri dishes 3 glass rods ultraviolet lamp 0.1 mol/l silver nitrate in dropper bottle 0.1 mol/l sodium chloride in dropper bottle 0.1 mol/l sodium bromide in dropper bottle 0.1 mol/l sodium iodide in dropper bottle Method 1. Ensure that your safety glasses are on. 2. Work in groups for this investigation. One group will investigate the effect of sunlight on the silver salts; a second group will investigate the effect of ultraviolet (UV) light on these salts; a third group will place the dishes in a dark cupboard. 3. Label three Petri dishes 1, 2 and 3. Put 20 drops of silver nitrate solution in each dish. 4. Add 20 drops of sodium chloride solution to dish 1 and mix with a stirring rod. Observe the reaction and record your observations. 5. Add 20 drops of sodium bromide solution to dish 2 and mix with a stirring rod. Observe the reaction and record your observations. 6. Add 20 drops of sodium iodide solution to dish 3 and mix with a stirring rod. Observe the reaction and record your observations. 7. One group will expose the precipitates in dishes 1, 2 and 3 to sunlight (e.g. near a window) for about 5 10 minutes. The second group will use the ultraviolet lamp to expose each dish in turn to UV light for the same amount of time. The third group will place the dishes in the dark as a control. Compare and record all results. 8. Clean up and return all equipment. CHAPTER 4 CHEMICAL EXTRACTION 87

15 DATA ANALYSIS Results and analysis 1. Construct a table to record your observations. Complete the table. 2. Compare the effects of sunlight and UV light on each silver salt explaining any differences. 3. Explain the importance of the control dishes. 4. Write balanced chemical equations for each precipitation reaction. 5. The effect of light on silver salts has a practical application in film (non-digital) photography. Photographic film is covered with a fine gelatine emulsion containing grains of silver bromide. The photosensitive material decomposes on exposure to light to form tiny clusters of silver atoms and bromine atoms at the surface of the crystalline grains. Exposure to more light produces more silver. The process is completed by the developer solution during film processing, resulting in a negative film in which the darkest areas correspond to areas where most metallic silver was formed. (a) A photon of light ejects an electron from a bromide ion to form a bromine atom. Write an ion/electron half-equation for this process. (b) Ejected electrons are accepted by silver ions, forming silver metal. Write an ion/electron half-equation for this process. Conclusion Write a brief conclusion for this experiment. 4.4 DATA ANALYSIS BOILING AND ELECTROLYSING WATER Part A: Boiling water Figure 4.14 shows a particle diagram of water heated on an electric hotplate to 100 C and 100 kpa. 1. Use the key in figure 4.14 to identify the particles present in the: (a) liquid state (b) vapour state. 2. Use figure 4.14 to explain why this process is a physical change rather than a chemical change. 3. Explain why the temperature of the water remains at 100 C throughout the process. Water molecule Hydrogen molecule Oxygen molecule Electric hotplate BOILING AND ELECTROLYSING WATER Water vapour Liquid water H 2 O (l) Heat Figure 4.14 Particle diagram of boiling water H 2 O (g) 88 THE CHEMICAL EARTH

16 DATA ANALYSIS Part B: Electrolysing water Figure 4.15 shows a particle diagram of the electrolysis of water at platinum electrodes in a voltameter after the electrolysis has been operating for some time. (A small amount of sodium sulfate was added to the water to make it conductive but this does not alter the results these ions are not shown.) 1. Use the key in figure 4.15 to identify the particles present in the water. 2. Use the key in figure 4.15 to identify the gases that have formed in the arms of the voltameter. 3. Use figure 4.15 to explain why this process is a chemical change rather than a physical change. 4. Suggest a reason why platinum is used for the electrode rather than some other metal. 5. The following half-equations describe processes occurring at the surface of each electrode. 2H 2 O(l) + 2e H 2 (g) + 2OH (aq) (1) 2H 2 O(l) O 2 (g) + 4H + (aq) + 4e (2) Which of these equations is consistent with the events occurring in figure 4.15 at the: (a) positive electrode? (b) negative electrode? 4.5 DATA ANALYSIS INVESTIGATING THE PURITY OF LIMESTONE BY THERMAL DECOMPOSITION The apparatus shown in figure 4.16 was used to determine the loss in weight as a limestone sample (impure calcium carbonate) was heated. Limestone does not decompose significantly over a Bunsen flame and so an electric furnace was used. The reaction tube was weighed and a sample of powdered limestone added to the tube. The tube and its contents were weighed and connected to the apparatus. Powdered sodium hydroxide was added to the U-tube to absorb the acidic carbon dioxide evolved on heating. The limestone was heated very strongly until no further decomposition took place. The reaction tube and its contents were reweighed after cooling to room temperature. Reaction tube Granules to absorb water vapour Gas Gas Water molecule Platinum electrodes Furnace Limestone or pure calcium carbonate Figure 4.16 Electric furnace decomposition of limestone and calcium carbonate Powdered sodium hydroxide absorbs carbon dioxide. + Hydrogen molecule Oxygen molecule Hydrogen ion Hydroxide ion + DC power supply Figure 4.15 Particle diagram of the electrolysis of water CHAPTER 4 CHEMICAL EXTRACTION 89

17 DATA ANALYSIS The experiment was repeated with a known sample of pure calcium carbonate rather than limestone. The results are shown in table 4.4. Table 4.4 Limestone Pure calcium carbonate Initial mass of reaction tube Initial mass of reaction tube + sample Final mass of reaction tube + sample g g g g g g 1. Calculate the percentage loss in weight on the thermal decomposition of the: (a) limestone (b) pure calcium carbonate. 2. Determine the percentage of pure calcium carbonate in the limestone. (Hint: Use your answers from question 1.) 3. State the assumptions that you have made in answering question For each experiment, determine the increase in weight of the U-tube. State the assumption that you are making. 5. The reaction in the U-tube can be represented by the word equation: carbon dioxide (gas) + sodium hydroxide (solid) sodium hydrogen carbonate (solid) Write a chemical equation for this reaction. 90 THE CHEMICAL EARTH