Chemistry IGCSE Textbook

NES/Chemistry/IGCSE Electroplating Electroplating is the process of depositing metals from solution in the form of a layer on other surfaces such as metal or plastic. The object to be plated is always the cathode. cathode The metal used to plate the object is the (object to anode. be plated) The electrolyte is always an aqueous salt electrolyte:- solution of the metal. If you are unsure of aqueous the solubility of a salt always use a nitrate copper(II) salt as all nitrates are soluble (see Topic nitrate 8.3) Cu(NO3)2 copper anode (metal used to plate the object) Anode At the anode copper atoms are oxidised, each losing two valence electrons forming copper(II) ions which enter the electrolyte, therefore the anode decreases in mass/size. The electrons travel around the external circuit to the cathode (key). Cathode At the cathode copper(II) ions in the electrolyte are reduced, each gaining two electrons forming copper atoms which are deposited on the surface of the clean key. The key has to be cleaned using steel wool or sand paper otherwise the deposited layer will not adhere to the surface. The key has to be continually rotated to uniformly coat the key with a layer of copper. 100

NES/Chemistry/IGCSE The longer the experiment is left running for, the thicker the layer of copper deposited. The closer the key is to the anode the higher the rate of reduction of copper(II) ions and so the faster the experiment proceeds (the bulb will glow brighter). If the object is to be The anode must be The electrolyte must be plated with made from copper copper aqueous copper(II) nitrate silver silver aqueous silver(I) nitrate zinc zinc aqueous zinc nitrate Uses of Electroplating  Silver is used to electroplate ornaments or cutlery, which has been made from copper or an alloy of copper. This makes it look more attractive and expensive.  Gold is used to electroplate ornamental items to make them look more expensive. The metal parts of electronic components are often gold plated to ensure the contacts will be free of corrosion and therefore provide good contacts. 101

NES/Chemistry/IGCSE Production of Electricity from Simple Cells A cell is a device which converts chemical energy into electrical energy and is composed of two metals of different reactivity connected by an external circuit and an electrolyte The bigger the difference in reactivity between the two metals, the higher the voltage, or the brighter the bulb.  Example 3: zinc - iron cell Zinc atoms, being more reactive than iron atoms, form their positive ions in preference to the iron. The zinc ions enter the electrolyte. Zn  Zn2+ + 2e– (oxidation) The electrons travel around the external circuit to the iron electrode. Here the hydrogen ions in the electrolyte remove the electrons from the iron electrode forming hydrogen molecules (H2). 2H+ + 2e–  H2 (reduction) voltmeter electron flow V positive negative terminal terminal less more reactive reactive metal (Fe) metal (Zn) bubbles of electrolyte hydrogen gas (aqueous sulphuric acid) The zinc anode decreases in mass/size but the iron cathode stays the same mass/size. 102

NES/Chemistry/IGCSE The voltage could be increased by replacing the Zn electrode with a piece of Mg metal or replacing the Fe electrode with a piece of copper metal. Use of Batteries Batteries are used as a convenient portable source of energy. Batteries are used to operate mobile telephones and personal stereos without the need to be connected to a mains supply of electricity. 103

NES/Chemistry/IGCSE Electrolysis of Concentrated Aqueous Sodium Chloride Solution (Brine) brine chlorine gas hydrogen gas Brine enters the electrolytic diaphragm cell. The ions present in solution are Na+, H+, Cl– and OH–. Na+ Na+ The diaphragm is used to H+ H+ Cl– OH– prevent the chlorine gas OH– H2O H2O produced from reacting aqueous with the sodium sodium hydroxide solution inside hydroxide the cell. power supply Material used for cathode - steel Material used for anode - titanium Anode At the anode chloride ions lose electrons and so are oxidised to form chlorine gas. 2Cl–(aq)  Cl2(g) + 2e– The cations now pass through the diaphragm but the Cl– ions cannot. Cathode At the cathode hydrogen cations gain electrons and are reduced to H2 molecules. 2H+(aq) + 2e–  H2(g) This leaves the ions Na+ and OH– ions which form aqueous sodium hydroxide. Overall Reaction 2NaCl(aq) + 2H2O(l)  2NaOH(aq) + Cl2(g) + H2(g) 104

NES/Chemistry/IGCSE Extraction of Aluminium by Electrolysis Aluminium cannot be extracted like iron and zinc because it is more reactive than carbon and cannot be reduced by carbon. So aluminium is made by the electrolysis of pure aluminium oxide also called alumina. Molten aluminium is produced at the cathode. The ore of aluminium is called bauxite. This contains aluminium oxide (Al2O3.2H2O) or alumina, which is amphoteric, and the impurity iron(III) oxide, which is basic. The ore is heated with aqueous sodium hydroxide. The aluminium oxide acts as an acidic oxide and reacts with the NaOH(aq) to form a solution of sodium aluminate. The basic Fe2O3 does not react with NaOH(aq). The mixture is then filtered to separate the insoluble Fe2O3 from the solution of sodium aluminate. Pure aluminium oxide is made from this solution. Solid pure aluminium oxide/alumina does not conduct electricity as there are no mobile electrons or mobile ions. The melting temperature of pure aluminium oxide is about 2050 oC, which is extremely high and so requires a lot of energy to be used. So molten cryolite (Na3AlF6) is used to dissolve the pure aluminium oxide/alumina which lowers the working temperature of the cell to about 900 oC, therefore the process is more efficient, saving a lot of energy. pure aluminium oxide dissolved in molten cryolite cathode anode (carbon (carbon in the form in the form of graphite) of graphite) molten outer casing aluminium 105

NES/Chemistry/IGCSE Also, the cryolite ionises producing many moles of ions, improving the electrical conductivity of the electrolyte and again the process is more efficient, saving energy. The electrolyte is kept molten due to the electrical heating effect of the current. The electrodes are made from the element carbon in the form of graphite. Cathode Molten aluminium is formed and can be siphoned/tapped off from the bottom. cathode reaction (reduction) Al3+(cryolite) + 3e–  Al(l) Anode Oxygen gas is formed at the anode, but F2(g), CO(g) and CO2(g) are also formed. anode reaction (oxidation) 2O2–(cryolite)  O2(g) + 4e– Fluorine from the cryolite is attracted to the anode and as the electrolysis is carried out at a high temperature, the oxygen formed at the anode reacts with the carbon anode forming CO(g) and CO2(g) which escape. This is why the anode decreases in mass and has to be replaced on a regular basis. Uses of Aluminium  Food containers because it is not poisonous, does not corrode and is malleable.  Aircraft because it has a low density and a high strength to weight ratio.  Electric power lines because it is a good conductor of electricity. Aluminium Oxide Aluminium is covered by a surface layer of unreactive, protective impermeable (to water and oxygen) aluminium oxide(Al2O3), which prevents (or slows) any further reactions with water, acids or heat. 106

NES/Chemistry/IGCSE Copper and Steel-Cored Aluminium Electrical Cables Copper electrical cable:  copper is a good conductor of electricity due to mobile valence electrons.  copper is ductile and so can be drawn into wires Steel cored aluminium cables: plastic aluminium  aluminium is used as it is a good coating cable conductor and has a low density steel  steel is used as it is a good conductor of core electricity and it is strong  This combination gives the greatest combination of strength, conductivity and low density. Recycling Recycling is where a material, such as an aluminium can, can be used again.  It conserves the reserves of the ore for the future.  It saves energy and fuels as large amounts of electricity are used to extract aluminium.  It saves money as extracting aluminium by electrolysis is an expensive process.  Recycling metal articles prevents them causing litter. Problems with Recycling Recycling is not always an economic possibility because of the cost of sorting, collecting and processing the waste material. The higher the value of the material, the more economical it is to recycle. Recycling plastics presents a particular problem because of the difficulty of identifying the type of polymer used. It is easy to separate iron from copper in recycling, but p.v.c. and poly(ethene) are much more difficult to distinguish and separate. 107

NES/Chemistry/IGCSE Topic 6 - Chemical Changes 6.1 Energetics of a reaction  Describe the meaning of exothermic and endothermic reactions  Interpret energy level diagrams showing exothermic and endothermic reactions  Describe bond breaking as an endothermic process and bond forming as an exothermic process  Draw and label energy level diagrams for exothermic and endothermic reactions using data provided  Calculate the energy of a reaction using bond energies 6.2 Energy transfer  Describe the release of heat energy by burning fuels  State the use of hydrogen as a fuel  Describe radioactive isotopes, such as 235U, as a source of energy  Describe the use of hydrogen as a fuel reacting with oxygen to generate electricity in a fuel cell. Details of the construction and operation of a fuel cell are not required. Definitions An exothermic reaction is one that releases heat energy into the surroundings An endothermic reaction is one which absorbs heat energy from the surroundings A fuel is a substance which can be conveniently used as a source of energy 108

NES/Chemistry/IGCSE 6.1 Energetics of a Reaction A change in energy, either an endothermic change or an exothermic change is a sign that a chemical reaction has taken place. Most chemical reactions are exothermic because reactants lose energy in the form of heat and light to form more energetically stable products of lower energy. In exothermic reactions the reactants are higher in energy than the products. In endothermic reactions the reactants are lower in energy than the products. Exothermic Reactions Endothermic Reactions Combustion Electrolysis Neutralisation Displacement Thermal decomposition Synthesis Photosynthesis Rusting Respiration Energy Level Diagrams The amount of energy stored in the reactants before the reaction is shown by the line on the left. The line on the right is the energy stored in the products. The difference between the two lines shows the energy given out during the reaction and is represented by the symbol H. Energy is measured in Joules (J) and kilojoules (kJ) so: 1kJ = 1000J Most reactions do not happen spontaneously as some energy is required to start the reaction by breaking the bonds of the reactants. This is called the activation energy (EA). 109

NES/Chemistry/IGCSE Exothermic Reactions Energy (kJ) reactants activation In an exothermic reaction, the H energy energy of the reactants is higher than the energy of the products products. The value for H will be negative. This is because the reactants are transferring (losing) energy to their surroundings. reaction coordinate Endothermic Reactions Energy (kJ) reactants activation In an endothermic reaction, energy the energy of the reactants is lower than the energy of the products products. The value for H will be positive. This is because H the reactants are transferring (gaining) energy from their surroundings. reaction coordinate Energy Changes when Bonds Break and Form Bond breaking is when bonds in reactants are broken, which requires energy and therefore is an endothermic process and H is positive. Bond forming is when bonds are made in a product, which releases energy and therefore is an exothermic process and H is negative. 110

NES/Chemistry/IGCSE  Example 1: The combustion of methane CH4 + 2O2  CO2 + 2H2O The intramolecular covalent bonds in the methane and oxygen molecules break. H is positive. The carbon forms two C=O bonds and four O–H bonds are also formed. H is negative. individual atoms 1 x C, 4 x O, 4 x H Energy (kJ) energy needed energy released when to break bonds new bonds are formed H OO O HH H C H+ H OO O C O+ O H is negative HH The first process involves the breaking of bonds which is endothermic and energy is absorbed. H is positive. The second process involves the forming of bonds which is exothermic and energy is released. H is negative. If more energy is released on forming bonds than is absorbed to break bonds, then the overall process is exothermic. If less energy is released on forming bonds than is absorbed to break bonds, then the overall process is endothermic. If we want to calculate the energy change in a reaction and know if it is exothermic, or endothermic, then we need to know bond energy values. 111

NES/Chemistry/IGCSE Bond Energy Bond energy is the amount of energy consumed or liberated when a bond is broken or formed in kJ/mol. It is not essential to learn the values in the table below, although you should be aware that different bonds require different amounts of energy to break them. Bond Energy needed to break bond Energy released on forming this bond (Endothermic) (kJ/mol) (Exothermic) (kJ/mol) H–H +436 -436 Cl–Cl +242 -242 H–Cl +431 -431 C–H +413 -413 C–C +347 -347 C–O +335 -335 O=O +497 -497 +740 -740 C=O +463 -463 O–H Step 1: How to Calculate the Change in Enthalpy Step 2: Step 3: Write out the equation Step 4: Draw out all the bonds in the reactants and the products Step 5: Calculate the energy required to break the reactant bonds Calculate the energy required to form the product bonds Calculate the enthalpy change using the equation H for the reaction = (bond breaking + bond forming) Remember - bond breaking values are positive - bond forming values are negative 112

NES/Chemistry/IGCSE  Example 2: calculating H for the combustion of methane Step 1 CH4 + 2O2  CO2 + 2H2O Step 2 H + OO OO HCH OO OCO + HH HH H Step 3 Bonds Broken (Endothermic) Type of Bond Heat Energy C–H x 4 kJ/mol +413 x 4 O=O x 2 +497 x 2 total amount of heat +2646 energy needed Step 4 Bonds Formed (Exothermic) Type of Bond Heat Energy kJ/mol C=O x 2 -740 x 2 O–H x 4 -463 x 4 total heat energy -3332 released Step 5 H for the reaction = (bond breaking + bond forming)      H = +2646 + (-3332) = -686 kJ/mol Therefore the reaction is exothermic. 113

NES/Chemistry/IGCSE 6.2 Energy Transfer Most fuels are combusted in oxygen to release the energy. A good fuel would:  be cheap  be available in large quantities  be a liquid at room temperature so is easily transported and stored  release a large amount of energy when combusted  not produce polluting gases. Combustion of Fuels In order for any material to combust (burn) three things must be present:  heat  a fuel  oxygen Combustion of Different Fuels - Comparing the Heat Energy Released thermometer  Use a measuring cylinder to put 100 copper can cm3 of water into the copper can spirit burner  Record the initial mass of the fuel on an electronic balance  Use a thermometer to record the initial temperature of the water  Light the fuel and allow to burn until the temperature rises by about 30oC  Record the final temperature of the water  Record the final mass of fuel  Repeat this for the other fuel  The best fuel produces the greatest temperature rise per gram of fuel 114

NES/Chemistry/IGCSE Hydrogen Advantages of Hydrogen Disadvantages of Hydrogen It releases large amounts of heat energy As hydrogen is a gas at room when it burns temperature, it cannot be easily stored and transported as it requires large Produces less pollution pressurised containers, which are very Renewable heavy. Hydrogen forms an explosive mixture with air making it very dangerous Uses of Hydrogen  reacting with oxygen to generate electricity in a fuel cell  to make ammonia  as a fuel for rockets  used to manufacture margarine from olive oil Hydrogen Fuel Cell –+ hydrogen oxygen hydrogen + oxygen water vapour porous anode containing a porous cathode containing a nickel catalyst nickel catalyst electrolyte (concentrated KOH(aq) 115

NES/Chemistry/IGCSE The hydrogen fuel cell is the most common type of fuel cell. It produces electricity from hydrogen and oxygen with the product being water. A hydrogen fuel cell is composed of a catalyst, anode, cathode and an electrolyte. A fuel cell produces electricity directly from the fuel. The electrons produced at the anode pass around the external circuit to the cathode. Anode Reaction H2(g) + 2OH–(aq)  2H2O(l) + 2e– Cathode Reaction O2(g) + 2H2O(l) + 4e–  4OH–(aq) Overall Reaction H2(g) + ½O2(g)  H2O(l) Uranium-235 Advantages of Uranium-235 Disadvantages of Uranium-235 A small amount of fuel releases large Radioactive, so needs careful handling amounts of heat energy and storage Does not burn, so does not produce Disposal of used uranium-235 fuel rods polluting gases Does not burn, so does not need oxygen 116

NES/Chemistry/IGCSE Topic 7 - Chemical Reactions 7.1 Physical and chemical changes  Identify physical and chemical changes, and understand the differences between them 7.2 Rate (speed) of reaction  Describe and explain the effect of concentration, particle size, catalysts (including enzymes) and temperature on the rate of reactions  Describe the application of the above factors to the danger of explosive combustion with fine powders (e.g. flour mills) and gases (e.g. methane in mines)  Demonstrate knowledge and understanding of a practical method for investigating the rate of a reaction involving gas evolution  Interpret data obtained from experiments concerned with rate of reaction  Devise and evaluate a suitable method for investigating the effect of a given variable on the rate of a reaction  Describe and explain the effects of temperature and concentration in terms of collisions between reacting particles (An increase in temperature causes an increase in collision rate and more of the colliding molecules have sufficient energy (activation energy) to react whereas an increase in concentration only causes an increase in collision rate.)  Describe and explain the role of light in photochemical reactions and the effect of light on the rate of these reactions (This should be linked to section 14.4.)  Describe the use of silver salts in photography as a process of reduction of silver ions to silver; band photosynthesis as the reaction between carbon dioxide and water in the presence of chlorophyll and sunlight (energy) to produce glucose and oxygen 117

NES/Chemistry/IGCSE 7.3 Reversible reactions  Understand that some chemical reactions can be reversed by changing the reaction conditions (Limited to the effects of heat and water on hydrated and anhydrous copper(II) sulfate and cobalt(II) chloride.) (Concept of equilibrium is not required.)  Predict the effect of changing the conditions (concentration, temperature and pressure) on other reversible reactions  Demonstrate knowledge and understanding of the concept of equilibrium 7.4 Redox  Define oxidation and reduction in terms of oxygen loss/gain. (Oxidation state limited to its use to name ions, e.g. iron(II), iron(III), copper(II), manganate(VII).)  Define redox in terms of electron transfer  Identify redox reactions by changes in oxidation state and by the colour changes involved when using acidified potassium manganate(VII), and potassium iodide. (Recall of equations involving KMnO4 is not required.)  Define oxidising agent as a substance which oxidises another substance during a redox reaction. Define reducing agent as a substance which reduces another substance during a redox reaction.  Identify oxidising agents and reducing agents from simple equations 118

NES/Chemistry/IGCSE 7.1 Physical and Chemical Changes  Melting point Physical Properties  Boiling point  Viscosity Temperature where a solid changes to a liquid  Brittle Temperature where a liquid forms vapor The ability of a liquid to flow  Malleable A material that breaks instead of bending when under  Ductile stress  Sonorous A material that can be bent, or hammered into shape  Solubility Can be stretched into a wire  Concentration Makes a ringing sound when hit  Density Ability of a substance to dissolve in a solvent  Electrostatic charge Amount of one substance in a mixture  Electrical conductivity Mass per unit volume of a substance  Thermal conductivity Positive, or negative charge on a particle A material's ability to conduct electricity Ability of a material to conduct heat Physical Changes These are changes that only involve a change in physical properties, such as  Melting Changing from solid to liquid  Freezing/Solidifying Changing from liquid to solid  Boiling/Evaporating Changing from liquid to gas  Condensing Changing from gas to liquid  Subliming Changing from solid to gas, or the opposite  Diffusion Spreading out, from an area of high concentration to low concentration by random motion. Physical changes are easy to reverse and do not involve new substances being made. 119

NES/Chemistry/IGCSE Chemical Properties  Valency Number of electrons that can be lost, gained, or shared during a chemical reaction  Forming Ions Forming positive ions (cations), or negative ions (anions)  Act as catalyst Speed up other chemical reactions, without being used up  Reactivity How reactive a chemical is  Forming coloured compounds  Forming acidic oxides  Forming basic oxide Chemical Changes Chemical changes occur when chemicals react together. New chemicals are made which will have new, different physical and chemical properties. Most chemical changes are difficult to reverse. All chemical changes can be written as word, or symbol, equations. Symbol equations must always be balanced. Word Equations Reagent A + Reagent B  Product C + Product D The chemicals at the start of a reaction are called reagents, they are always on the left hand side of the equation. The new chemicals at the end of a reaction are called products, they are always on the right hand side of the equation. The + sign means \"with\" The  sign means \"changes into\"  Example 1: When iron reacts with sulphur a new chemical called iron sulphide is made. This new substance will have different chemical and physical properties to iron and sulphur. iron + sulphur  iron sulphide 120

NES/Chemistry/IGCSE Symbol Equations These equations show the symbols and formula of the reagents and products. Symbol equations must always be balanced, so the number of each type of atom is the same on the left hand side as the right hand side.  Example 2: Write a symbol equation for the reaction of iron and sulphur, producing iron (II) sulphide. Fe + S  FeS  Example 3: Write a symbol equation for the reaction of calcium hydroxide with hydrochloric acid, producing calcium chloride and water. Ca(OH)2 + 2 HCl  CaCl2 + 2H2O 121

NES/Chemistry/IGCSE 7.2 Rate (Speed) of Reaction Definitions Activation energy is the minimum energy required to break the bonds of the reagent particles Rate of a chemical reaction is the concentration of reagent used up, or product made, in a given time A catalyst is a substance which speeds up a chemical reaction, but remains chemically unchanged at the end Enzymes are biological catalysts which speed up reactions, but remain chemically unchanged at the end A photochemical reaction is one where light causes a reaction to occur. The higher the light intensity the higher the rate of the reaction Collision Theory In order for a reaction to occur there must be successful collisions between reagent particles. This requires: 1. The reacting particles (reagents) must collide 2. The particles must collide with enough energy - called activation energy, EA. If particles do not collide, or do not have enough energy when they do collide, then no reaction takes place and the particles keep moving. Rate of Reaction The higher the number of successful collisions per unit time, the higher the rate of reaction. Most reactions start of at a maximum rate of reaction and then slow down. This is because the concentration of reagents is the greatest just as the reaction starts, then as the reagents get used up, their concentration decreases. Once the reaction is over the concentration of the reagents is usually zero. 122

NES/Chemistry/IGCSE Units of Rate of Reaction This is the change of concentration, per unit time. Concentration is measured in mol/dm3 and time is usually measured in s. The rate of a reaction is measured in mol/dm3/s Factors Affecting Rate of Reaction 1. Concentration (of a solution) 2. Temperature 3. Particle Size 4. Pressure (of a gas) 5. Catalyst, or Enzyme 1. Concentration of Solutions Increasing concentration of reagents increases the rate of reaction. This is because there are more particles per unit volume, so the collision rate between reacting particles increases, therefore the successful collision rate increases, which results in an increased rate of reaction. lower concentration higher concentration 2. Temperature Increasing temperature increases the rate of reaction. This is because the average kinetic energy of the particles increases, which means they are moving faster. So more particles have an energy greater than, or equal to the activation energy, therefore the collision rate and the successful collision rate increases; resulting in an increased rate of reaction. For every 100C the temperature increases, the rate of reaction doubles (approximately). 123

NES/Chemistry/IGCSE 3. Particle Size Decreasing the particle size (increasing surface area) increases the rate of reaction. This is because there are more reagent particles exposed to collide, so the collision rate increases. The successful collision rate increases, resulting in an increased rate of reaction. larger particle size smaller particle size (smaller surface area) (larger surface area) Very small particles (powder) can be dangerous as they will react very quickly. Any type of dry food, or solid fuel powder may easily explode if ignited (set on fire). 4. Pressure of Gases Increasing the pressure in a gaseous system increases the rate of reaction. The distance between particles is reduced as the pressure increases. There are more particles per unit volume, so the collision rate increases, therefore the successful collision rate increases, resulting in an increased rate of reaction. Low pressure High pressure Less particles per unit volume More particles per unit volume 124

NES/Chemistry/IGCSE 5a. Catalyst Adding a catalyst increases the reaction rate. A catalyst allows the reaction to go by an alternative pathway with a lower activation energy. More particles will have an energy greater than, or equal to the activation energy on collision, therefore the successful collision rate increases, resulting in an increased reaction rate. uncatalysed energy (kJ) reagents catalysed products Increasing the surface area of a catalyst will also increase the rate of the reaction. There can often be more than 1 catalyst for a particular reaction.  Example 1: The decomposition of hydrogen peroxide 2H2O2(l)  2H2O(l) + O2(g) Without a catalyst this reaction takes a long time With manganese(IV) oxide (black powder), the reaction takes only seconds. When a piece of liver (contains iron(III) ions) is added the reaction takes only seconds. 5b. Enzymes Enzymes are proteins, complex organic molecules (we will study them in more detail in Topic 14). Most enzymes only work in a narrow range of temperature, or pH. 125

NES/Chemistry/IGCSE Measuring Rates of Reaction This can be done by:  Measuring volume of gas made, per unit time  Measuring the mass loss, per unit time  Measuring the colour change, per unit time  Measuring the increase in turbidity (cloudiness of a suspension), per unit time Measuring Volume of Gas Made Reagents that will produce a gas on mixing must be chosen. The gas is collected and measured; usually in a gas syringe, or an inverted measuring cylinder. The reagents are best kept separate so as to avoid losing any gas at the start. The experiment can be repeated varying concentration of solution, temperature, particle size, pressure of a gas, or a catalyst. All other variables must be kept the same to ensure a fair test for the comparison of results. 126

NES/Chemistry/IGCSE  Example 2: mixing calcium carbonate with hydrochloric acid CaCO3(s) + 2HCl(aq)  CaCl2(aq) + H2O(l) + CO2(g) gas syringe suspended test tube containing CaCO3(s) HCl(aq) stop clock  Using an electronic balance weigh out a specific mass of calcium carbonate and place inside the suspended test tube  Using a measuring cylinder add a specific volume (excess) of dilute hydrochloric acid into the conical flask.  Set the gas syringe to read zero  Place the bung in the conical flask  Shake the conical flask (to tip over the test tube allowing the acid and calcium carbonate to react) and start the stop clock simultaneously  Record the volume of gas produced every ten seconds for five minutes 127

volume of gas (cm3) NES/Chemistry/IGCSE Results - Changing Particle Size Plotting the graph of volume (y-axis) against time (x-axis) would produce the following graphs, using different particles sizes as a comparison. powdered fine coarse time (s) Dilute hydrochloric acid reacts vigorously with powdered calcium carbonate and the reaction is completed in the shortest time. This is because the powder has a smaller particle size (larger surface area) and so more reagent particles are exposed to collide, so the collision rate is higher than in the fine and coarse experiments, therefore the successful collision rate is higher, resulting in a higher rate of reaction The initial rate of reaction is usually greatest because there is the highest concentration of acid particles present and also the greatest mass of CaCO3(s) present at the start, so the collision rate is the highest, therefore the successful collision rate is the highest, resulting in a the highest reaction rate. The graphs level off (volume of gas remains constant) showing that the CaCO3(s) has been completely used up (the limiting reagent). If all the graphs level off at the same final volume, then a fair test has been carried out. The reaction rate decreases as time increases because the concentration of HCl(aq) decreases with time and the mass of CaCO3(s) decreases, so the collision rate decreases, therefore the successful collision rate decreases, resulting in a decreasing reaction rate. 128

Volume of gas (cm3) NES/Chemistry/IGCSE volume of gas (cm3)Results - Changing Solution Concentration Plotting the graph of volume (y-axis) against time (x-axis) would produce the following graphs, using different acid concentrations as a comparison. 3 mol/dm3 2 mol/dm3 1 mol/dm3 CaCO3 is the limiting reagent time (s) Limiting Reagent is CaCO3 (fixed variable) - so changing acid concentration has no effect on volume of gas made. excess CaCO3 used 3 mol/dm3 2 mol/dm3 1 mol/dm3 time (s) Limiting Reagent is HCl (independent variable) - so changing acid concentration will affect the volume of gas made. For both graphs, the initial rate will increase with increased concentration of acid. 129

NES/Chemistry/IGCSE  Example 3: The decomposition of hydrogen peroxide 2H2O2(aq)  2H2O(l) + O2(g) This reaction happens very slowly at room temperature and pressure. The rate of reaction increases considerably if a catalyst is used. The volume of oxygen gas made per unit time can be used to measure the rate of the reaction. Different catalysts will having varying effects on the rate of reaction.  Place a specific volume of hydrogen peroxide into the conical flask  Add a specific mass of manganese(IV) oxide (catalyst) into the suspended test tube and place in the conical flask  Put the bung into position  Use a gas syringe to record the volume of gas produced every 10 seconds The experiment can be repeated using different catalysts to see which speeds up the reaction the most. Other variables (mass of catalyst, particle size of catalyst, volume of hydrogen peroxide and temperature) must be kept the same to make a fair test for the comparison of results. 130

NES/Chemistry/IGCSE Measuring the Mass Loss This method can be used when there is a loss in mass due to one of the products being gaseous. An electronic balance is used to measure the mass change during the experiment. The experiment can be repeated varying concentration of solution, temperature, particle size, pressure of a gas, or a catalyst. All other variables must be kept the same to ensure a fair test for the comparison of results.  Example 4: mixing calcium carbonate with hydrochloric acid CaCO3(s) + 2HCl(aq)  CaCl2(aq) + H2O(l) + CO2(g) cotton wool plug tare -0.11 electronic balance  Place a specific mass of powdered calcium carbonate in a conical flask and place on the electronic balance  Measure out a specific volume (excess) of hydrochloric acid and place on the electronic balance  Press the “tare” to set the balance reading to zero.  Add the hydrochloric acid to the calcium carbonate and insert the cotton wool plug.  Record the mass every ten seconds. A cotton wool plug is used to allow the CO2(g) to escape, but stop acid spray from effervescence. 131

NES/Chemistry/IGCSEloss in mass (g) Results Obtained from the Experiment Plotting the graph of mass loss (y-axis) against time (x-axis) would produce the following graphs, using different particle size as a comparison. powdered fine coarse time (s) All graphs level off at the same mass loss. This shows that equal masses of calcium carbonate have been used for each experiment. Measuring the Increase in Turbidity (Cloudiness)  Example 5: Sodium thiosulfate reacting with hydrochloric acid, producing a precipitate of sulfur. Na2S2O3(aq) + 2HCl(aq)  2NaCl(aq) + H2O(l) + SO2(g) + S(s) yellow solid cross drawn  As the reaction takes place more and on paper more precipitate of sulfur is made.  This makes the flask cloudier.  After some time, the cross is no longer visible from above.  To ensure a fair comparison the depth of liquid in the flask must be kept constant. 132

concentration mol/dm3 NES/Chemistry/IGCSE 1/time (s–1) Add aqueous sodium thiosulfate into a conical flask  Add hydrochloric acid to the conical flask  Swirl the flask to mix and start the stop clock simultaneously  Place the conical flask on the cross drawn on a piece of paper  Stop the stop clock when the solution turns cloudy and the cross cannot be seen anymore Results - Changing Solution Concentration time (s) The graph shows as concentration increases, time taken for the reaction decreases, so rate of reaction increases. 0 concentration (mol/dm3) Rate (1/time) of reaction is directly proportional to concentration. If you double the concentration you double the rate of reaction and the time taken halves. 133

temperature (oC) NES/Chemistry/IGCSE 1/time (s–1). Results - Changing Temperature Plot a graph of temperature against time and then plot a graph of 1/time (rate) against temperature. time (s) temperature (oC) As temperature increases, rate of reaction increases. However, rate is not directly proportional to temperature. As a general rule the rate doubles for every 100C rise in temperature. 134

NES/Chemistry/IGCSE Explosive Reactions Any reaction that happens very quickly is dangerous as it is explosive.  Any dry substance that can burn may explode if it is a fine powder. This includes flour powder/dust in flour mills and coal powder/dust in coal mines.  Gases can build up in poorly ventilated rooms and explode. This includes methane gas in coal mines. Refrigerators  Food will decompose (go off) when left for a few days. This is a chemical reaction and can be slowed down by reducing the temperature. Most refrigerators keep food at 30C and this makes the food last much longer. Photochemical Reactions Light is a form of energy and provides energy for some reactions. It is not a catalyst. There are only a few chemical reactions that are affected by light: 1. Film photography - silver(I) ion reduction 2. Photosynthesis 3. Halogenation of alkanes 1. Silver(I) Ions The silver halides (AgCl, AgBr, AgI) are all photo sensitive (sensitive to light). This is how older film photography works. On a piece of photographic film the chemical is cream silver(I) bromide crystals. Light causes the silver(I) ions in the silver(I) bromide crystals to gain electrons and are reduced to silver atoms. Ionic half-equation Ag+ + e–  Ag . Br–  Br2 + 2e– Ionic half-equation Symbol equation 2AgBr(s)  2Ag(s) + Br2(g) The more light that shines on the photographic film, the darker it gets. So the rate of reaction is proportional to the light intensity. 135

NES/Chemistry/IGCSE 2. Photosynthesis Chlorophyll molecules absorb light by a process called photosynthesis. 6CO2(g) + 6H2O(l)  C6H12O6(aq) + 6O2(g) 3. Halogenation of Alkanes Alkanes react with halogens, in the presence of UV light to form halogenoalkanes. CH4 + Cl2  CH3Cl + HCl 136

NES/Chemistry/IGCSE 7.3 Reversible Reactions A reversible reaction is a reaction in which reagents form products and the product(s) can then react or decompose to form the reagents Dynamic equilibrium is reached in a closed system when the rate of the forward reaction and the rate of the backward reaction are equal and the concentration of reagents and products remain constant Reagents react together to form products. A+BC+D In some reactions, when the products are made, they can react together to form the reagents. These are called reversible reactions. C+DA+B Instead of writing two separate equations, they are combined into one equation. A+B⇌C+D The ⇌ symbol means a reversible reaction.  Example 1: hydrated copper(II) sulfate hydrated copper(II) sulfate ⇌ anhydrous copper(II) sulfate + water ⇌CuSO4.5H2O(aq) CuSO4(s) + 5H2O(l) Reading the equation forwards: when the hydrated copper(II) sulfate is heated, it decomposes into anhydrous copper(II) sulfate and steam. The process is endothermic and requires heating, which is why the water is steam. Reading the equation backwards: when water is added to anhydrous copper(II) sulfate it forms hydrated copper(II) sulfate. The process is exothermic. 137

NES/Chemistry/IGCSE Equilibrium This is when a reaction is reversible and the forward and backward reactions happen at the same time. All the reagents and products will be present in the equilibrium mixture. At dynamic equilibrium:  Rate of forward reaction = rate of backward reaction  The concentrations of reagents and products are constant A+B⇌C+D Concentration at start maximum minimum (zero) As the reaction happens decreases increases Rate of reaction at start maximum minimum (zero) As the reaction happens decreases increases At some point, the rate of the forward reaction must be equal to the rate of the backward reaction. This is called dynamic equilibrium. Once the rates of reaction are equal, the concentrations will not change any more and become constant. rate of reaction rate of forward reaction decreases with time dynamic equilibrium rate of backward reaction increases with time time (s) 138

NES/Chemistry/IGCSE Even though the forward and backward reactions are still occurring, there is no visible change observed because the rate of the two reactions in opposite directions are equal. Dynamic equilibrium can only be achieved in a closed system (materials cannot enter or leave the system) like a sealed gas jar. Position of Equilibrium The amount of reagent and product in an equilibrium reaction can vary and does not have to be 50% reagent and 50% product. If a reaction has more than 50% reagent, we say that the equilibrium position lies to the left and sometimes the symbol is used. If a reaction has more than 50% product, we say that the equilibrium position lies to the right and sometimes the symbol is used.  Example 2: The Haber Process (for making ammonia) N2(g) + 3H2(g) ⇌ 2NH3(g) The equilibrium position lies to the left. In fact the %Yield (how much product there is) is 15%. This means that there is 85% reagent.  Example 3: The Contact Process (for making sulfuric acid) 2SO2(g) + O2(g) ⇌ 2SO3(g) The equilibrium position lies to the right. In fact the %Yield is 90%. This means there is 90% product and 10% reagent. 139

NES/Chemistry/IGCSE Changing Equilibrium Position The equilibrium position can be moved:  To the left - making more reagent at dynamic equilibrium  To the right - making more product at dynamic equilibrium We can alter the equilibrium position by altering the: 1. Concentration of reagents or products 2. Temperature of the system 3. Pressure (if there are gases as reagents, or products) 4. Adding/removing reagent/product 5. Adding acid/alkali to a reaction with H+/OH- ions Le Chatelier's Principle If you increase the concentration of reagents or products, or increase the temperature, or pressure in a sealed reaction vessel; the equilibrium position shifts in such a way as to try to cancel what you are doing. In other words, the equilibrium position will shift in such a way as to do the opposite of what changes away from room temperature and pressure we do. 1. Changing Concentration If we increase the concentration of a reagent the equilibrium position will shift to the right as the system tries to use up the extra reagent that we added. If we reduce the concentration of a reagent then the opposite will happen. If we increase the concentration of a product the equilibrium position will shift to the left as the system tries to use up the extra product that we added. If we reduce the concentration of a product then the opposite will happen. 140

NES/Chemistry/IGCSE What are we Equilibrium Change in Concentration Concentration doing to the Position will Equilibrium of reagents of products shift to try to: will will system? Position increasing decrease decrease increase concentration concentration of moves from of reagents left to right increase decrease decreasing reagents concentration increase moves from increase decrease of reagents concentration of right to left increasing reagents decrease increase concentration decrease moves from of products concentration of right to left decreasing products concentration increase moves from of products concentration of left to right products 2. Changing Temperature Equilibrium reactions are exothermic in one direction and endothermic in the opposite direction. Increasing the temperature will favour the endothermic reaction. Decreasing the temperature will favour the exothermic side. What are we doing to Equilibrium Position Moves in the exothermic/ the system? will shift to try to endothermic direction endothermic direction increasing temperature decrease the temperature of system of the system exothermic direction decreasing temperature increase temperature of of system the system  Example 4: The Haber Process N2 + 3H2 ⇌ 2NH3 H = -92kJ The forward direction has an enthalpy change that is negative, so it is exothermic. Thus the backward reaction will have an enthalpy change that is positive and be endothermic. Increasing the temperature will cause the equilibrium position to shift to the left towards the endothermic side of the equilibrium reaction. 141

NES/Chemistry/IGCSE Decreasing the temperature will cause the equilibrium position to shift to the right towards the exothermic side of the equilibrium reaction.  Example 5: Reacting a carboxylic acid with an alcohol (esterification) CH3COOH + C2H5OH ⇌ CH3COOC2H5 + H2O H = 0 kJ Changing the temperature will have no effect on this reaction as there is no change in enthalpy. 3. Changing Pressure If there are gases in the reaction then changing the pressure will affect the equilibrium position. Increasing the pressure will favour the side with less moles of gas. Decreasing the pressure will favour the side with more moles of gas. What are we doing to Equilibrium Position Shifts to the side with more mole of gas or the system? shifts to try to: less mole of gas increasing the pressure decrease the pressure less mole of gas decreasing the pressure increase the pressure more mole of gas  Example 6: Thermal decomposition of calcium carbonate CaCO3(s) ⇌ CaO(s) + CO2(g) The left hand side has 0 moles of gas The right hand side has 1 mole of gas Increasing the pressure will shift the equilibrium position to the left. Decreasing the pressure will shift the equilibrium position to the right. 142

NES/Chemistry/IGCSE  Example 7: Formation of hydrogen iodide H2(g) + I2(g) ⇌ 2HI(g) The left hand side has 2 moles of gas The right hand side has 2 moles of gas Changing the pressure will have no effect on the equilibrium position as the number of moles left and right are equal. 4. Adding / Removing some Reagent / Product  Adding reagent - this will cause the equilibrium position to move to the right hand side.  Removing reagent - this will cause the equilibrium position to move to the left hand side.  Adding product - this will cause the equilibrium position to move to the left hand side.  Removing product - this will cause the equilibrium position to move to the right hand side. In each case the equilibrium position will shift to try to negate the changes we do to the system. Adding Acid / Alkali to a Reaction with H+ / OH- Ions Equilibrium reactions with H+ ions: Adding acid will increase the concentration of H+ ions and the equilibrium position will shift to the side without H+ ions. Adding alkali will decrease the concentration of H+ ions and the equilibrium position will shift to the side with H+ ions. 143

NES/Chemistry/IGCSE Equilibrium reactions with OH- ions: Adding acid will decrease the concentration of OH- ions and the equilibrium position will shift to the side with the OH- ions. Adding alkali will increase the concentration of OH- ions and the equilibrium position will shift to the side without the OH- ions.  Example 8: Bromine reacting with water Br2(aq) + H2O(l) ⇌ Br–(aq) + 2H+(aq) + OBr–(aq) orange colourless Adding acid will shift the equilibrium position to the left. Adding alkali will shift the equilibrium position to the right. You would see the following colour changes: intense orange colourless orange H+ OH– What are we Equilibrium position From left to Concn of Concn of right or right to doing? shifts to try to: reagents will products will left? increasing H+ decrease H+ concn right to left increase decrease concn increase H+ concn left to right decrease increase decreasing H+ concn by adding OH– 144

NES/Chemistry/IGCSE 5. Adding a Catalyst A catalyst does not change the equilibrium position, but the equilibrium reaction will reach dynamic equilibrium faster. A catalyst increases the rate of both the forward and backward reaction by equal amounts. Different metals, or metal compounds, affect different reactions. So you need to get the right catalyst for a specific reaction. Some reactions are affected by more than one catalyst, and each catalyst will have a different affect on the rate of reaction. Enzymes are biological catalysts and affect specific reactions only. For example protease affects the rate of decomposition of protein. 145

NES/Chemistry/IGCSE 7.4 Redox Oxidation is loss of electrons Reduction is gain of electrons A redox reaction is a reaction in which one species has been oxidised and another species has been reduced Oxidising agent is a substance which oxidises another substance during a redox reaction Reducing agent is a substance which reduces another substance during a redox reaction Oxidation means:  loss of electrons  gain of oxygen  loss of hydrogen Reduction means:  gain of electrons  loss of oxygen  gain of hydrogen Oxidising agents Reducing agents oxygen hydrogen Group VII (halogens) Group VII ions (halides) acidified potassium dichromate(VI) Group I metals acidified potassium manganate(VII) magnesium hydrogen peroxide carbon monoxide methane gas 146

NES/Chemistry/IGCSE Oxidation Number Oxidation number is the charge on a species. It is the same as the valencies given in Topic 4.1. Working out the Oxidation Numbers  An atom is neutral and will always have a charge of zero  A molecule is neutral and will always have a charge of zero  The charges of the atoms in a molecule, or compound will add up to zero  The charges of the atoms in a polyatomic ion add up to the charge of the polyatomic ion  Oxygen has a charge of -2 in compounds, except H2O2 where it has a charge of -1 and OF2 where it has a charge of +2  Fluorine always has a charge of -1 in a compound  Metal ions are always positive  Example 1: carbon dioxide, CO2 The molecule has an overall charge of zero Each oxide ion has a charge of -2 So C + (-2 x 2) = 0 C = +4  Example 2: phosphorous pentachloride, P2O5 The molecule has an overall charge of zero Each oxide ion has a charge of -2 So 2P + (-2 x 5) = 0 2P = +10 P = +5 147

NES/Chemistry/IGCSE  Example 3: nitrate ion, NO31- The polyatomic ion has a charge of -1 Each oxide ion has a charge of -2 So N + (-2 x 3) = -1 N = -1 + 6 N = +5  Example 4: chromate(VI) ion, Cr2O72- The polyatomic ion has a charge of -2 Each oxide ion has a charge of -2 So 2Cr + (-2 x 7) = -2 2Cr = -2 + 14 2Cr = +12 Cr = +6 Working out the Change in Oxidation Numbers The change in oxidation number can be worked out if the oxidation number of a species is know at the start (as a reagent) and end (as a product) of a reaction.  Example 5: Mg(s) + Cu(NO3)2(aq)  Mg(NO3)2(aq) + Cu(s) Charges of magnesium Mg = 0 Mg in Mg(NO3)2 = +2 Change in oxidation number = +2 Magnesium has been oxidised Cu(NO3)2 is acting as an oxidising agent Charges of copper Cu in Cu(NO3)2 = +2 Cu = 0 Change in oxidation number = -2 Copper has been reduced Mg is acting as a reducing agent 148

NES/Chemistry/IGCSE Colour Changes in Redox Reactions Acidified potassium manganate(VII) is an oxidising agent. It changes from purple to colourless when oxidising other chemicals. The manganate(VII) ion is itself is reduced to a Mn2+ ion. Acidified potassium chromate(VI) is an oxidising agent. It changes from orange to green when oxidising other chemicals. The chromate(VI) ion is itself is reduced to a Cr3+ ion. Strength of Reducing Agents Very reactive metals are strong reducing agents, because they can readily lose their valence electrons to another species. The more reactive a metal is, the stronger a reducing agent it will be. A more reactive metal can reduce the ions of a less reactive metal, displacing it from a solution of its salt or compound.  Example 6: metal displacement reaction Symbol equation 3Mg + Fe2(SO4)3  3MgSO4 + 2Fe Ionic equation 3Mg + 2Fe3+  3Mg2+ + 2Fe Ionic half-equations 3Mg  3Mg2+ + 6e– oxidation reaction 2Fe3+ + 6e–  2Fe reduction reaction So magnesium metal can reduce iron(III) cations to form iron atoms 149