4.9 describe experiments to carry out acid-alkali titrations

• Have a known volume of acid in a beaker with methyl orange, the solution will be red as it is very acidic.
• Set up a burette with alkali in it. Open the tap very slightly, so that it drips very slowly into the acid.
• With each drop stir the contents of the beaker.
• The more alkali that is added the more neutral and closer to orange it gets.
• When the solution is neutral the it will be completely orange, at this point close the tap on the burette.
• The level in the burette will have dropped, showing the volume of alkali used.
• This shows you how much alkali you need to use to neutralise the acid.

4.8 describe experiments to prepare insoluble salts using precipitation reactions

Silver nitrate and sodium chloride are added together, the product, silver chloride is made, this salt is insoluable and so will form a white precipitate in the solution.

AgNO3 + NaCl > AgCl + NaNO3

4.7 describe experiments to prepare soluble salts from acids

Dilute sulphuric acid is added to an excess of magnesium.
Mg + H2SO4 > MgSO4 + H2
The left over magnesium is filtered off and the mixture boiled down slowly to concentrate.
When it is cooled, crystals will form, these can be blotted dry with a piece of paper.

This can be done with other metals and acids:

Nitric, sulphuric and hydrochloric acids make soluble salts with most metals.

ammonium, potassium and sodium make soluble salts with acids

4.25 predict the effects of changing the pressure and temperature on the equilibrium position in reversible reactions.

If you move the equilibrium, you change the rate of reaction.
If the equilibrium moves to the right, you have more products. (the reactants are reacting faster)
If the equilibrium moves to the left, you have more reactants. (the products are reacting faster)

If you increase the pressure: the equilibrium will move to the side with least molecules.
If you decrease the pressure: the equilibrium will move to the side with the most molecules.

If you increase the temperature, there will be more products that are produced by an endothermic reaction.
This is because the reaction is trying to use up the extra heat, and it does so by putting the energy into making bonds.

4.24 understand the concept of dynamic equilibrium

Dynamic equilibrium is when a reversible reaction is happening both ways at the same time, at the same rate.

4.23 describe reversible reactions such as the dehydration of hydrated copper(II) sulfate and the effect of heat on ammonium chloride

If you add water to copper sulphate you can make hydrated copper sulphate.
If you remover the water from hydrated copper sulphate you can make copper sulphate.

When heated, ammonium chloride splits into to hydrogen chloride and ammonia.
Hydrogen chloride and ammonia can be reacted to make ammonium chloride.

4.22 understand that some reactions are reversible and are indicated by the symbol ⇌ in equations

Some reactions can happen both ways: the reactants can make the products and the products can make the reactants.

This symbol shows that ⇌.

4.21 explain that a catalyst speeds up a reaction by providing an alternative pathway with lower activation energy

A catalyst provides an alternative route for the reaction to start, this route requires less energy to start the reaction.

4.18 describe the effects of changes in surface area of a solid, concentration of solutions, pressure of gases, temperature and the use of a catalyst on the rate of a reaction

Higher temperature, more surface area, higher concentration, higher pressure and use of a catalyst all make a reaction faster.

4.19 understand the term activation energy and represent it on a reaction profile

Activation energy is the amount of energy required for a reaction start happening.

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4.20 explain the effects of changes in surface area of a solid, concentration of solutions, pressure of gases and temperature on the rate of a reaction in terms of particle collision theory

Collision theory says that to react particles must:
Collide with enough energy to react
Collide in the right orientation to react (the more frequent the collisions, the more likely this is)

Surface area

• Particles collide more frequently if there is more surface area, as there is more contact between the reactants. Faster rate of reaction.

Concentration/ pressure

• There is more chance of particles colliding at a higher concentration/pressure, so they react more often. Faster rate of reaction.

Temperature

• Particles move about more and will collide more frequently the higher the temperature; react more often. Increases the rate of reaction.

Catalyst

• Provides an alternative pathway for the reaction to start which requires a lower activation energy.

4.17 describe experiments to investigate the effects of changes in surface area of a solid, concentration of solutions, temperature and the use of a catalyst on the rate of a reaction

Surface area

• Put a set mass of magnesium in hydrochloric acid
• Time the reaction
• Change the from of magnesium keeping the mass the same (powder, wire, strips)
• The more surface area (the smaller the pieces of magnesium) the faster the reaction

Concentration

• Put a set mass of marble chips into dilute hydrochloric acid
• Time the reaction
• Change the ratio of water to hydrochloric acid
• The more concentrated the hydrochloric acid (the lower the ratio of water) the faster the reaction

Temperature

• Put a set mass of magnesium powder into a set mass of hydrochloric acid
• time the reaction
• Carry out this reaction at different temperatures
• The higher the temperature the faster the rate of reaction

Catalyst

• If you have hydrogen peroxide it will not decompose
• If you put it with manganese dioxide it will decompose into water and oxygen
• The manganese dioxide will be unaltered by the reaction
• The more of the catalyst the faster the reaction

4.16 use average bond energies to calculate the enthalpy change during a simple chemical reaction.

To work out the enthalpy change you do:
 the sum of bond energies in the reactants - the sum of bond energies in the products

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))

Then we substitute in:

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

Credit to author of this source: https://www.google.co.uk/url?sa=t&rct=j&q=&esrc=s&source=web&cd=5&cad=rja&ved=0CFoQFjAE&url=http%3A%2F%2Fchemistryatdulwich.wikispaces.com%2Ffile%2Fview%2FSection%2B4%2Bb%2Benergetics.doc%2F183320535%2FSection%25204%2520b%2520energetics.doc&ei=a72UUZOaHYa_0QXMkoD4Bg&usg=AFQjCNF9yn2bAWyGOPYSO4W1NrJExx7HQw&sig2=_XxB8HEYh8MHVUe5abc6sA&bvm=bv.46471029,d.d2k

4.15 understand that the breaking of bonds is endothermic and that the making of bonds is exothermic

Breaking bonds requires energy, this takes in heat: endothermic.

Making bonds releases energy, this gives out heat: exothermic.

4.14 represent exothermic and endothermic reactions on a simple energy level diagram

Exothermic: lower energy level at end;

gcsescience

Endothermic: more energy at end;

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Activation energy is the energy requires to start the reaction.

4.13 understand the use of ΔH to represent enthalpy change for exothermic and endothermic reactions

ΔH is the symbol that represents the amount of energy lost or gained in a reaction.

+ΔH is endothermic (because it gains heat)
-ΔH is exothermic (because it looses heat)

4.12 calculate molar enthalpy change from heat energy change

ΔH =	T (C) x mass of H2O (g) x 4.2J/g C./
	number of moles

ΔH is measured in J/mol or kJ/mol.

• ΔH = molar enthalpy change
• T = temperature increase or decrease caused by the reaction
• mass of H₂O = mass of water in grams;
• 4.2J/g C = specific heat capacity of water
  
  
  Credit to who ever made this source: https://www.google.co.uk/url?sa=t&rct=j&q=&esrc=s&source=web&cd=5&cad=rja&ved=0CFoQFjAE&url=http%3A%2F%2Fchemistryatdulwich.wikispaces.com%2Ffile%2Fview%2FSection%2B4%2Bb%2Benergetics.doc%2F183320535%2FSection%25204%2520b%2520energetics.doc&ei=a72UUZOaHYa_0QXMkoD4Bg&usg=AFQjCNF9yn2bAWyGOPYSO4W1NrJExx7HQw&sig2=_XxB8HEYh8MHVUe5abc6sA&bvm=bv.46471029,d.d2k

4.11 describe simple calorimetry experiments for reactions such as combustion, displacement, dissolving and neutralisation in which heat energy changes can be calculated from measured temperature changes

Measure the temperature at the beginning of the experiment, measure the temperature at the end. How ever much heat has gone up or down is the calorimetry of the reaction.

For example if you have a beaker of water and take its temperature, then burn a piece of bread under it, the change in temperature is the calories (energy) of the bread.

4.10 understand that chemical reactions in which heat energy is given out are described as exothermic and those in which heat energy is taken in are endothermic

In an exothermic reaction heat is given out. Because bonds are made which gives out energy.
Think expelled; exit; exo.

In a endothermic reaction heat is taken in. Because bonds are Broken which requires energy.

4.6 understand the general rules for predicting the solubility of salts in water

i) all common sodium, potassium and ammonium salts are soluble

ii) all nitrates are soluble

iii) common chlorides are soluble, except silver chloride

iv) common sulfates are soluble, except those of barium and calcium

v) common carbonates are insoluble, except those of sodium, potassium and ammonium