# GCSE & A Level Revision Notes

### Acids, Alkalis and Salts

#### 4.1

Litmus – natural indicator

• Red litmus turns blue in alkalis
• Blue litmus turns red in acids

Phenolphthalein – Pink in alkalis – Clear in acids

Methyl Orange – Red in strong acids (below around 4)

• Yellow in weak acids and alkalis

Universal Indicator – gives a quantitative result for how acidic or alkaline something is. Is red in acids and blue in alkalis.

#### 4.2

The pH scale, which ranges from 0 to 14, can classify solutions as strongly acidic, weakly acidic, neutral, weakly alkaline or strongly alkaline.

#### 4.3

Universal indicator can be used to test for the approximate pH value of different solutions.

#### 4.4

Acids can be defined as sources of H+ ions whereas alkalis can be defined as a source of OH- ions.

#### 4.5

Dilute acid + metal → salt + hydrogen

• Hydrochloric acid: metal chloride salt
• Nitric acid: metal nitrate salt
• Sulphuric acid: metal sulphate salt

Metal oxide + acid → salt + water Metal hydroxide + acid → salt + water Metal Carbonate + acid → salt + carbon dioxide + water

#### 4.6

General Salt solubility rules

• All sodium, potassium and ammonium salts are soluble
• All nitrates are soluble
• Common chlorides are soluble except for silver chloride
• Common sulphates are soluble except those of barium and calcium
• Common carbonates are insoluble except sodium, potassium and ammonium.
• Common hydroxides are insoluble except those of sodium, potassium and ammonium.
 SOLUBILITY RULES

#### 4.7

To prepare soluble salts that are not those of ammonium, potassium or sodium, you can react acid and the metal together, which will produce the salt and hydrogen. You can then get a pure sample of the salt by evaporating off most of the liquid until crystals start to appear and then allowing the rest to evaporate naturally from an evaporating dish, left in a cool, dry place.

#### 4.8

To prepare an insoluble salt, two soluble salts are added together and the insoluble salt made precipitates out, which is called a precipitation reaction. For example, if silver nitrate and sodium chloride solutions are added together, which will produce silver chloride (insoluble) and sodium nitrate. The mixture is then filtered, leaving behind only the insoluble salt, which is then washed will distilled water and left to dry to give a pure sample of the insoluble salt.

#### 4.9

To prepare salts from potassium, sodium and ammonium, you must use an acid-alkali titration. The acid should be in the burette, with the alkali in the beaker/conical flask. Put phenolphthalein into the beaker, which will indicate when the solution has gone neutral. Do a test run to see roughly how much acid is needed to neutralise the alkali. When the indicator changes from pink to clear, the solution is neutralised and a salt will be produced. To get a pure sample of the salt, do the titration without any indicator once you establish exactly how much acid is needed.

### Energetics

#### 4.10

Chemical reactions can either give out energy in the form of heat or take it in. Those that give out energy are known as exothermic reactions whereas those that take it in are known as endothermic reactions.

#### 4.11

To test for the molar enthalpy change when combusting fuel, this apparatus is used. The mass of the water and the temperature of the water are taken before the experiment as well as the mass of fuel before burning. After the experiment, the temperature is recorded again, as well the mass of the fuel afterwards.

 You use the equation shown in 4.12 to calculate the enthalpy change in total and then the mass of the fuel burned to calculate it per mole. A simple calorimeter is used to investigate reactions of displacement, dissolving and neutralisation. The temperature of the solution before and after the experiment is recorded, as well as both the concentration and volume of the solution, or the mass of the solution so that the enthalpy change can be calculated per mole.

#### 4.12

ΔH is the molar enthalpy change ΔT is the change in temperature Mass of H2O is the mass of the water in grams 4.2 J/g ˚C is the specific heat capacity of water

#### 4.13

ΔH is used to represent the enthalpy change of a reaction. In an exothermic reaction, where energy is given out ΔH is negative, whereas in an endothermic reaction, ΔH is positive.

#### 4.15

In chemical reactions, breaking bonds requires energy, which means that it is endothermic, whereas making bonds releases energy and is therefore exothermic.

#### 4.16

To work out the enthalpy change in a reaction you subtract the sum of the bond energies in the products from the sum of the bond energies in the reactants.

Given average bond energies, you can easily work out the sum of bond energies. For example, if we are told that C-H = 413 O=O= 498 C=O= 746 H-O=464

Work out the enthalpy change for this reaction: CH4 + 2O2 → CO2 + 2H2O

(4 x C-H) + (2 x O=O) - (2x C=O) + (4 x H-O)

((4 x 413) + (2 x 498)) – ((2 x 746) + (4 x 464)) = 2648KJ - 3348KJ ΔH = -700KJ

### Rates of Reaction

#### 4.17

Rate of reaction is affected by a number of different factors, including surface area, concentration, temperature and use of a catalyst. You can investigate the effects of these factors using simple experiments:

### Surface Area

• Use marble chips in different sizes: large, medium, small, powdered
• Put the same mass of marble chips into hydrochloric acid (25cm3)
• Time the reaction
• You should observe the higher the surface area, the quicker the rate of reaction.

### Concentration

• Use different ratios of acid and water: 25:75, 50:50, 75:25
• Use a strip of magnesium ribbon the same length each time
• Time the reaction
• You should observe that the higher the concentration of acid, the quicker the rate of reaction.

### Temperature

• Do the reaction at different temperatures: 10˚C, 20˚C, 30˚C, 40˚C 50˚C
• Use a strip of magnesium ribbon the same length each time
• Time the reaction
• You should observe that the higher the temperature, the faster the rate of reaction.

### Catalyst

• Hydrogen peroxide decomposes very slowly naturally
• If you add manganese dioxide it will decompose much faster into water and oxygen
• The manganese dioxide will be unaltered as it is a catalyst
• You should observe that the more catalyst there is, the faster the rate of reaction will be.

#### 4.18

The higher the surface area, temperature and concentration, the faster the rate of reaction will be, the same if you add more of a catalyst.

#### 4.19

The activation energy is the energy required for a reaction to start. It is represented on an energy diagram as the spike between the reactants and products. It can be written as EA.

 Reaction Progress Graph

#### 4.20

Particle collision theory says that particles must collide with enough energy to react in order for a reaction to take place. The greater the surface area, the more area there is for particles to react with, and the higher the concentration, the more particles there are to react. If the temperature is increased, the particles have more kinetic energy, which means they are more likely to collide and have enough energy to react.

#### 4.21

A catalyst speeds up a reaction by providing an alternative pathway for the reaction, with lower activation energy.

### Equilibria

#### 4.22

Some reactions are reversible, which means they go both ways. In equations, an arrow that points in both directions (⇌) represents this.

#### 4.23

An example of a reversible reaction is the dehydration of hydrous copper sulphate. When heated, the water of crystallisation evaporates, leaving anhydrous white crystals. If you re-add water to the crystals, it will be reabsorbed, giving you hydrous copper sulphate again. Ammonium chloride splits into ammonia and chlorine when heated. If you heat it in a closed system, when it reaches the top, which is cooler, it will reform into ammonium chloride.

#### 4.24

Dynamic equilibrium occurs when the forward and backward reaction are taking place at an equal and constant rate.

#### 4.25

You can move the equilibrium to change the amount of reactants or products being made in the reaction. One way to do this is by increasing or decreasing temperature and pressure. If you increase the pressure, equilibrium will move to the side with the least moles of substance, because it will take up less space, and if you decrease the pressure, it will move to the side with more moles, as this allows the substance space to expand. If you increase the temperature, the endothermic side of the reaction will be favoured whereas if you decrease the temperature then the exothermic side of the reaction will be favoured. However, when the temperature is very low, there is a very poor rate of reaction and so not much reactant or product is made. This means that chemists have to find a balance between equilibrium and rate of reaction.