# AQA Chemistry 2b

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## Rate of Reaction

1) One of the slowest is the rusting of iron.

2) A moderate speed reaction is a metal (like magnesium) reacting with an acid to producea gentle stream of bubbles.

3) A really fast reaction is an explosion, where its all over in a fraction of a second.

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## Depends on Four Things

1) Temperature

2) Concentration (or pressure for gases)

3) Catalyst

4) Surface area of solids (or size of solid pieces)

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## Typical Graphs

The quickest reaction is shown by the line with the steepest slope.

The faster a reaction goes, the sooner it finishes, which means the line becomes flat earlier.

Graph:

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## Explanations

1) Graph 1 represents the original fairly slow reaction. The graph is not too steep.

2) Graphs 2 and 3 represent the reaction taking place quicker but with the same initial amounts. The slope of the graph gets steeper.

3) The increased rate could be due to any of these:

increase in temperature, increase in concentration (or pressure), catalyst added, solid reactant crushed up into smaller bits

4) Graph 4 produces more product as well as going faster. This can only happen if more reactant(s) are added at the start. Graphs 1, 2 and 3 all converge at the same level, showing that they all produce the same amount of product, although they take different times to get there.

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## Measuring the Rate

The rate of reaction can be observed either by measuring how quickly the reactants are used up or how quickly the products are formed. It's usually a loteasier to measure the products forming. The rate of reaction can be calculated using the following formula:

Rate of Reaction = Amount of reactant used or amount of product formed / Time

There are different ways that a rate of reaction can be meassured.

Learn these three:

• Precipitation
• Change in Mass (Usually Gas Given Off)
• The Volume of Gas Given Off
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## Precipitation

1) This is when the product of the reaction is a precipitate which clouds the solution.

2) Observe a mark through the solution and measure how long it takes to disappear.

3) The quicker the mark disappears, the quicker the reaction.

4) This only works for reactions where the initial solution is rather see-through.

5) This result is very subjective - different people may not agree over the exact point when the mark 'disappears'.

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## Change in Mass

1) Measuring the speed of a reaction that produces a gas can be carried out on a mass balance.

2) As the gas is released the mass disappearing is easily measure on the balance.

3) The quicker the reading on the balance drops, the faster the reaction.

4) Rate of reaction graphs are particularly easy to plot using the results from this method.

5) This is the most accurate of the three methods described because the mass balance is very accurate. But it has the disadvantage of releasing the gas straight into the room.

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## Volume of Gas Given Off

1) This involves the use of a gas syringe tp measure the volume of gas given off.

2) The more gas given off during a given time interval, the faster the reaction.

3) A graph of gas volume against time elapsed could be plotted to give a rate of reaction graph.

4) Gas syringes usually give volumes accurate to the nearest millimetre, so they're quite accurate. You have to be quite careful though - if the reaction is too vigorous, you can easily blow the plunger out of the end of the syringe.

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The mass balance method is only accurate as long as the flask isn't too hot, otherwise you lose mass by evaporation as well as by the reaction.

The first method isn't very accurate, but if you're not producing a gas you can't use either of the other two.

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## Hydrochloric Acid and Marble Chips

This experiment is often used to demonstrate the effect of breaking the solid up into smaller bits.

1) Measure the volume of gas evolved with a gas syringe and take readings at regular intervals.

2) Make a table or readings and plot them as a graph. You choose regular time intervals, and time goes on the x-axis and volume goes on the y-axis.

3) Repeat the experiment with exactly the same volume of acid, and exactly the same volume of marble chips, but with the marble more crunched up.

4) Then repeat with the same mass of powdered chalk instead of marble chips.

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## Effect of Using Finer Particles

1) Using finer particles means that the marble has a larger surface area.

2) A larger surface area causes more frequent collisions so the rate of reaction is faster.

3) Line 4 shows the reaction if a greater mass of small chips is added. The extra surface area gives a quicker reaction and there is also more gas evolved overall.

Graph:

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## Magnesium Metal with Dilute HCL

1) This reaction is good for measuring the effects of increased concentration (as is the marble/acid reaction).

2) This reaction gives off hydrogen gas, which we can measure with a mass balance.

3) In this experiment, time also goes on the x-axis and volume goes on the y-axis.  (The other method is to use a gas syringe.)

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## More Concentrated Acid Solutions

1) Take readings of mass at regular time intervals. Put the results in a table and work out the loss in mass for each reading. Plot a graph.

2) Repeat with more concentrated acid solutions, but always with the same amount of magnesium. The volume of acid must always be kept the same too - only the concentration is increased.

3) The three graphs show the same pattern - a higher concentration giving a steeper graph, with the reaction finishing much quicker.

Graph:

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## Sodium Thiosulfate and HCL

1) These two chemicals are both clear solutions. They react together to form a yellow precipitate of sulfur. The experiment involves watching a black mark through the cloudy sulfur and timing how long it takes to go.

2) The reaction can be repeated for solutions at different temperatures. This is hard to do accurately and safely - use a water bath to heat the solutions to the right temperature before you mix them. The depth of the liquid must be kept the same each time.

3) The results of course show that the higher the temperature the quicker the reaction and therefore the less time it takes the mark to disappear. Typical results:

Temperature (degrees C):          20           25          30         35         40

Time taken to disappear (s):     193         151         112       87         52

This reaction can also be used to test the effects of concentration. This reaction does not give a set of graphs.

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## Decomposition of Hydrogen Peroxide

This is a good reaction for showing the effect of different catalysts. The decomposition of hydrogen peroxide is: 2H2O2 --> 2H2O + O2

1) This is normally quite slow but a sprinkle of maganese(IV) oxide catalyst speeds it up no end. Other catalyst that work are found in a) potato peel and b) blood.

2) Oygen gas is given off, which provides an ideal way to measure the rate of reaction using a gas syringe.

3) Same graphs. Better catalysts give a quicker reaction, which is shown by a steeper graph which levels off quickly. This reaction can also be used to measure the effects of temperature, or of concentation of the H2O2 solution. The graphs will look the same.

Graph:

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## Collision Theory

Reaction rates are explained by collision theory. The rate of a reaction depends on how often and how hard the reacting particles collide with each other. The basic idea is that the particles have to collide in order to react, and they have to collide hard enough (with enough energy).

More collisions increases the rate of reaction.

The effects of temperature, concentration and surface area on the rate of reaction can be explained in terms of how often the reacting particles collide successfully.

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## More Collisions

Higher temperature increases collisions:

• When the temperature is increased the particles all move quicker. If they're moving quicker, they're going to collide more often.

Higher concentration (or pressure) increases collisions:

• If a solution is made more concentrated it means that there are more particles of reactant knocking about between the water molecules which makes the collisions between the important particles more likely. In a gas, increasing the pressure means the particles are more squashed together so there will be more frequent collisions.

Larger surface area increases collisions:

• If one of the reactants is a solid then breaking it up into smaller pieces will increase the total surface area. This means the particles around it in the solution will have more area to work on, so there'll be more frequent collisions.
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## Faster Collisions

Faster collisions increase the rate of reaction.

Higher temperature also increases the energy of the collisions, because it makes all the particles move faster.

Increasing the temperature causes faster collisions.

Reactions only happen if the particles collide with enough energy.

The minimum amount of energy needed by the particles to react is known as the activation energy.

At a higher temperature there will be more particles colliding with enough energy to make the reaction happen.

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## Catalysts

Catalysts speed up reactions.

Many reactions can be speeded up by adding a catalyst.

A catalyst is a substance which speeds up a reaction, without being changed or used up in the reaction.

A solid catalyst works by giving the reacting particles a surface to stick to.

This increases the number of successful collisions (and so speeds up the reaction).

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## Reduce Costs in Industrial Reactions

1) Catalysts are very important for commercial reasons - most industrial reactions use them.

2) Catalysts increase the rate of the reaction, which saves a lot of money simply because the plant doesn't need to operate for as long to produce the same amount of stuff.

3) Alternatively, a catalyst will allow the reaction to work at a much lower temperature. That reduces the energy used up in the reaction (the energy cost), which is good for sustainable development and can save a lot of money too.

4) They can be very expensive to buy, and often need to be removed from the product and cleaned. They never get used up in a reaction though, so once you've got them you can use them over and over again.

5) Different reactions use different catalysts, so if you make more than one product at your plant, you'll probably need to buy different catalysts for them.

6) Catalysts can be 'poisoned' by impurities, so they stop working, e.g. sulfur impurities can poison the iron catalyst used in the Haber process (used to make ammonia for fertilisers). That means that you have to keep your reaction mixture very clean.

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## Exothermic Reaction

An exothermic reaction is once which transfers energy to the surroundings, usually in the form of heat and usually shown by a rise in temperature.

1) The best example of an exotheric reaction is buring fuels - also called combustion. This gives out a lot of heat - it's very exothermic.

2) Neutralisation reactions (acid+alkali) are also exothermic.

3) Many oxidation reactions are exothermic. For example, adding sodium to water produces heat, so it must be exothermic. The sodium emits heat and moves about on the surface of the water as it is oxidised.

4) Exothermic reactions have lots of everyday uses. For example, some hand warmers use the exothermic oxidation of iron in air (with a salt solution catalyst) to generate heat. Self heating cans of hot chocolate and coffee also rely on exothermic reactions between chemicals in their bases.

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## Endothermic Reaction

An endothermicreaction is one which takes energy in from the surroundings, usually in the form of heat and is usually shown by a fall in temperature.

Endothermic reactions are much less comon. Thermal decompositions are a good example:

Heat must be supplied to make calcium carbonate decompose to make quicklime.

CaCO3 --> CaO + CO2

Endothermic reactions also have everyday uses. For example, some sports injury packs use endothermic reactions - they take heat in and the pack becomes very cold. More convenient than carrying ice around.

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## Reversible Reactions

In reversible reactions, if the reaction is endothermic in one direction, it will be exothermic in the other direction. The energy absorbed by the endothermic reaction is equal to the energy released during the exothermic reaction.

A good example is the thermal decomposition of hydrated copper sulfate.

endothermic

hydrated copper sulfate                         anhydrous copper sulfate + water

exothermic

("anhydrous" just means without water and "hydrated" means with water)

1) If you heat blue hydrated copper(II) sulfate crystals it drives the water off and leaves white anhydrous copper(II) sulfate powder. This is endothermic.

2) If you then add a couple of drops of water to the white powder you get the blue crystals back again. This is exothermic.

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## pH Scale

• The pH scale is a measure of how acidic or alkaline a solution is.
• The strongest acid has pH 0. The strongest alkaline has pH 14.
• A neautral substance has pH 7 (e.g. pure water).

pH 1 - car battery acid, stomach adic

pH 3 - vinegar, lemon juice

pH 4 - acid rain

pH 5 - normal rain

pH 8/9 - washing-up liquid

pH 10 - pancreatic juice

pH 11 - soap powder

pH 12 - bleach

pH 13/14 - caustic soda (drain cleaner)

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## Indicator

• An indicator is just a dye that changes colour.
• The dye in the indicator changes colour depending on whether it's above or below a certain pH.
• Universal indicator is a combination of dyes which gives the colours on the pH scale.
• It's very useful for estimating the pH of a solution.
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## Acids & Bases

Acids and bases neutralise each other.

• An acid is a substance with a pH less than 7. Acids form H+ ions in water.
• A base is a substance with a pH greater than 7.
• An alkali is a base that dissolves in water. Alkalis for OH- ions in water.
• So, H+ ions make solutions acidic and OH- ions make them alkaline.

The reaction between acids and bases is called neutralisation: acid + base --> salt + water

Neutralisation can also be seen in terms of H+ and OH- ions: H+(aq) + OH-(aq) --> H2O(I)

Hydrogen (H+) ions react with hydroxide (OH-) ions to produce water.

When an acid neutralises a base (or vice versa), the products are neutral, i.e. they have a pH of 7. An indicator can be used to show that a neutralisation reaction is over (Universal indicator will go green).

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## State Symbols

State symbols tell you the physical state.

(s) - Solid

(l) - Liquid

(g) - Gas

(aq) - Aqueous (dissolved in water)

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## State Symbols

State symbols tell you the physical state.

(s) - Solid

(l) - Liquid

(g) - Gas

(aq) - Aqueous (dissolved in water)

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## Metals React With Acids

Acid + Metal --> Salt + Hydrogen

1) The more reactive a metal, the faster the reaction will go - very reactive metals (e.g. sodium) react explosively.

2) Copper does not react with dilute acids at all - because it's less reactive than hydrogen.

3) The speed of reaction is indicated by the rate at which the bubbles of hydrogen are given off.

4) The hydrogen is confirmed by the burning splint test - 'squeaky pop'.

5) The name of the salt produced depends on which metal is used and which acid is used.

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## Which salt is produced?

Hydrochloric acid will always produce chloride salts:

2HCl + Mg --> MgCl2 + H2 (Magnesium chloride)

Sulfuric acid will always produce sulfate salts:

3H2SO4 + 2Al --> Al2(SO4)3 + 3H2 (Aluminium sulfate)

Nitric acid produces nitrate salt when NEUTRALISED, but...

Nitric acid reacts fine with alkalis, to produce nitrates, but it can react with metals and produce nitrogen oxides instead.

Metals that are less reactive than hydrogen don't react with acid. Some metals (like sodium and potassium) are too reactive to mix with acid in a school lab - the beaker would explode.

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## Metal Oxides & Metal Hydroxides

Metal oxides and metal hydroxides are bases.

Some metal oxides and metal hydroxides dissolve in water. These soluable compounds are alkalines.

Even bases that won't dissolve in water will still react with acids,

All metal oxides and metal hydroxides react with acids to form a salt and water.

Acid + Metal Oxide --> Salt + Water

Acid + Metal Hydroxide --> Salt + Water

These are neutralisation reactions.

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## Deciding the Salt

hydrochloric acid + copper oxide --> copper cloride + water

sulfuric acid + calcium hydroxide --> calcium sulfate + water

nitric acid + potassium hydroxide --> potassium nitrate + water

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## Neutralising Ammonia

Ammonia can be neutralised with HNO3 to make fertiliser.

Ammonia dissolves in water to make an alkaline solution.

When it reacts with nitric acid, you get a neutral salt - ammonium nitrate:

NH3 (aq) + HNO3 (aq) --> HN4NO3 (aq)

Ammonia + Nitric acid --> Ammonium nitrate

This is a bit different from most neutralisation reactions because there's NO WATER produced - just the ammonium salt.

Amonium nitrate is an especially good fertiliser because it has nitrogen from two sources, the ammonia and the nitric acid. Plants need nitrogen to make proteins.

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## Making Salts

• If you're making a salt it's important to know if it's soluable or not so you know which method to use.
• Most chlorides, sulfates and nitrates are soluable in water.

• Most oxides and hydroxides are insoluable in water.
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## Soluble Salts: Metal or Insoluble Base

1) You need to pick the right acid, plus a metal or an insoluble base (a metal oxide or metal hydroxide). E.g. if you want to make copper chloride, mix hydrochloric acid and copper oxide.

CuO (s) + 2HCl (aq) --> CuCl2 (aq) + H2O (l)

2) You add the metal, metal oxide or hydroxide to the acid - the solid will dissolve in the acid as it reacts. You know when all the acid has been neutralised because the excess solid will just sink to the bottom of the flask.

3) Then filter out the excess metal, metal oxide or hydroxide to get the salt solution. To get pure, solid crystals of the salt, evaporate some of the water (to make the solution more concentrated) and then leave the rest to evaporate very slowly. This is called crystallisation.

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## Soluble Salts: Alkali

1) You can't use the other method with alkalis (soluble bases) like sodium, potassium or ammonium hydroxides, because you can't tell when the reaction has finished - you can't just add an excess to the acid and filter out what's left.

2) You have to add exactly the right amount of alkali to neutralise the acid - you need to use an indicator to show when the reaction's finished. Then repeat using exactly the same volumes of alkali and acid so the salt isn't contaminated with indicator.

3) Then just evaporate off the water to crystallise the salt as normal.

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## Making Insoluble Salts

1) If the salt you want to make is insoluble, you can use a precipitation reaction.

2) You need to pick two solutions that contain the ions you need. E.g. to make lead chloride you need a solution that contains lead ions and one which contains chloride ions. So you can mix lead nitrate solution (most nitrates are soluble) with sodium chloride solution (all group 1 compounds are soluble).

E.g.           Pb(NO3)2 (aq) + 2NaCl (aq) --> PbCl2 (s) + 2NaNO3 (aq)

3) Once the salt has precipitated out (and is lying at the bottom of your flask), all you have to do is filter it from the solution, wash it and then dry it on filter paper.

4) Precipitation reactions can be used to remove poisonous ions (e.g. lead) from drinking water. Calcium and magnesium ions can also be removed from water this way - the make water "hard", which stops soap lathering properly. Another use of precipitation is in treating effluent (sewage) - again, unwanted ions can be removed.

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## Electrolysis

Electrolysis means splitting up with electricity.

1) If you pass an electrical current through an ionic substance that's molten or in solution, it breaks down into the elements it's made of. This is called electrolysis.

2) It requires a liquid to conduct the electricity, called the electrolyte.

3) Electrolytes contain free ions - they're usually the molten or dissolved ionic substances.

4) In either case it's the free ions which conduct the electricity and allow the whole thing to work.

5) For an electrical circuit to be complete, there's got to be a flow of electrons. Electrons are taken away from ions at the positive electrode and given to other ions at the negative electrode. As ions gain or lose electrons they become atoms or molecules and are released.

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## Oxidation and Reduction

Core chemistry: reduction = loss of oxygen. However, reduction is also a gain of electrons.

On the other hand, oxidation is a gain of oxygen or loss of electrons.

So, "reduction" and "oxidation" don't always have to involve oxygen.

Electolysis ALWAYS involves an oxidation and a reduction.

il         eduction

s            s

oss      ain

Remember it as OIL RIG.

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## Electrolysis of Molten Lead Bromide

When a salt (e.g. lead bromide) is molten it will conduct electricity.

1) +ve ions are attracted to the -ve electrode. Here they gain electrons (reduction).

Lead is produced at the -ve electrode.

At the -ve electrode, one lead ion accepts two electons and becomes one lead atom.

2) -ve ions are attracted to the +ve electrode. Here they lose electrons (oxidation).

Bromine is produced at the +ve electrode.

At the +ve electrode, two bromide ions lose one electron each and become one bromine molecule.

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## Reactivity Affects the Products

1) Sometimes are more than two free ions in the electrolyte.

For example, if a salt is dissolved in water there will also be some H+ and OH- ions.

2) At the negative electrode, if metal ions and H+ ions are present, the metal ions will stay in solution if the metal is more reactive than hydrogen. This is because the more reactive an element, the keener it is to stay as ions. So, hydrogen will be produced unless the metal is less reactive than it.

3) At the positive electrode, if OH- and halide ions (Cl-, Br-, I-) are present then molecules of chlorine, bromine or iodine will be formed. If no halide is present, then oxygen will be formed.

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## Electrolysis of Sodium Chloride Solution

When common salt (sodium chloride) is dissolved in water and electrolysed, it produces three useful products - hydrogen, chlorine and sodium hydroxide.

1) +ve ions are attracted to the -ve electrode. Here they gain electons (reduction).

Hydrogen is produced at the -ve electrode. (H+ ions are released from the water.)

At the negative electrode, two hydrogen ions accept two electrons to become one hydrogen molecule.

2) -ve ions are attracted to the +ve electrode. Here they lose electrons (oxidation).

Chlorine is produced at the +ve electrode.

At the positive electrode, two chloride (Cl-) ions lose their electrons and become one chlorine molecule.

3) The sodium ions stay in solution because they're more reactive than hydrogen. Hydroxide ions from water are also left behind. This means that sodium hydroxide (NaOH) is left in the solution.

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## Half Equations

Half equations show the reactions at the electrodes.

The main this is to make sure the number of electons is the same for both half equations.

(You need to make sure the atoms are balanced too.)

For the electrolysis of sodium chloride the half-equations are:

Negative electrode:           2H+ + 2e- --> H2

Positive electrode:            2Cl- --> Cl2 + 2e-       OR       2Cl- - 2e- --> Cl2

For the electrolysis of molten lead bromide the half-equations are:

Negative electrode:           Pb2+ + 2e- --> Pb

Positive electrode:            2Br- --> Br2 + 2e-

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## Useful Products

The products of the electrolysis of sodium chloride solution are pretty useful in industry.

1) Chlorine has many uses, e.g. in the production of bleach and plastics.

2) Sodium hydroxide is a very strong alkali and is used widely in the chemical industry, e.g. to make soap.

3) Hydrogen can be used as a fuel or in the manufacture of margarine.

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## Removing Aluminium from its Ore

1) Aluminium is a very abundant metal, but it is always found naturally in compounds.

2) Its main ore is bauxite, and after mining and purifying, a white powder is left.

3) This is pure aluminium oxide, Al2O3.

4) The aluminium has to be extracted from this using electrolysis.

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## Cryolite

Cryolite is used to lower the temperature (and costs).

1) Al2O3 has a very high melting point of over 2000 degrees C - so melting it would be very expensive.

2) Instead the aluminium oxide is dissolved in molten cryolite (a less common ore of aluminium.)

3) This brings the temperature down to about 900 degrees C, which makes it much cheaper and easier.

4) The electrodes are made of carbon (graphite), a good conductor of electricity.

5) Aluminium forms at the negative elecrode and oxygen forms at the positive electrode.

Negative Electrode: Al3+ + 3e- ---> Al          Positive Electrode: 2O2- ---> O2 + 4e-

6) The oxygen then reacts with the carbon in the electrode to produce carbon dioxide. This means that the positive electrodes gradually get 'eaten away' and have to be replaced every now and again.

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## Electroplating Uses Electrolysis

1) Electroplating uses electrolysis to coat the surface of one metal with another metal, e.g. you might want to electroplate silver onto a brass cup to make it look nice.

2) The negative electrode is the metal object you want to plate and the positive electrode is the pure metal you want to plate it with. You also need an electrolyte to contain ions of the plating metal. (The ions that plate the metal object come from the solution, while the positive electrode keeps the solution 'topped up'.)

EXAMPLE: To electroplate silver onto a brass cup, you'd make the brass cup the negative electrode (to attract positive silver ions), a lump of pure silver the positive electrode and dip them into a solution of silver ions, e.g. silver nitrate.

3) There are lots of different uses for electroplating:

• Decoration: Silver is attractive, but very expensive. It's much cheaper to plate a boring brass cup with silver, than it is to make the cup out of solid silver - but it looks just as pretty.
• Conduction: Metals like copper conduct electricity well - because of this they're often used to plate metals for electronic circuits and computers.
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