Chemistry 4 - Chemical Changes

Question 1

Acids and bases

Answer 1

The pH scale goes from 0 to 14
The lower the pH of a solution the more acidic it is
The higher the pH of a solution the more alkaline it is
A neutral substance has pH 7

Measure the ph of a solution

• an indicator ( a dye that changes colour depending on whether it above or below a certain pH)
• pH probe ( used to measure pH electronically, attached to a pH meter, displayed as a numerical number)

Acids and bases neutralise each other
- a base is a substance with a pH greater than 7

Acid + base > salt + water

Question 2

Titrations

Answer 2

Allow you to find out exactly how much acid is needed to neutralise a quantity of alkaline

1) using a pipette and pipette filler, add a set volume of the alkali to a conical flask. Add two or three drops of indicator too
2) use a funnel to fill a burette with some acid of known concentration ( below eye level for safety) record initial volume of acid in burette
3) using the burette, add the acid to the alkaline a bit at a time - giving the conical flask a regular swirl. Go especially slowly when you the end-point (colour change is about to be reached)
4) the indicator changes colour when all the alkaline has been neutralised
5) record the final volume of the acid in the burette and use it along with the initial volume, to calculate the volume of acid used to neutralise the alkali

Use single indicators for titrations ( so you can see a sudden colour change)

• litmus ( blue in alkalis- red in acids)
• phenolphthalein ( pink in alkalis - colourless in acids)
• methyl orange ( yellow in alkalis - red in acids)

Question 3

Strong and Weak acids

Answer 3

Acids produce protons in water
- they ionise in aqueous solutions ( they produce h+ ions - hydrogen ions)

Strong acids ionise completely in water e.g sulfuric acid

Weak acids only partly ionise in water e.g ethanoic or citric acid

Ionisation of a weak acid is a reversible reaction ( only a few h+ ions released so the equilibrium lies well to the left

Strong acids are more reactive as if the concentration of h+ ions is higher the rate of reaction will be faster

pH is a measure of the concentration of h+ ions in a solution
For every decrease of 1 on the pH scale the concentration of h+ ions increases by 10

Factor H+ ion concentration changes by =
10^-x
If pH falls from 7 to 4 then x = -3

Question 4

Reactions of acids

Answer 4

Metal oxides and metal hydroxides are bases
Acid + Metal Oxide > Salt + water
Acid +Metal Hydroxide > salt + water

Acids and Metal Carbonates produce Carbon Dioxide
Acid + Metal Carbonate > Salt + Water + Carbon Dioxide e.g
Hydrochloride acid + sodium carbonate > sodium chloride + water + Carbon dioxide

Can make soluble salt using an insoluble base

1) gently warm the dilute acid using a Bunsen burner, then turn it off
2) add the insoluble base to the acid a bit at a time until no more reacts ( the base is in excess - the excess solid will sink to the bottom)
3) then filter out the excess solid to get the salt solution
4) to then get pure solid Salt crystals you can use crystallisation

Question 5

Reactivity series

Answer 5

Lists metals in order of their reactivity towards other substances 
Reactivity decreases as you go down this list :
Potassium - K
Sodium - Na
Lithium - Li
Calcium - Ca
Magnesium - Mg
Carbon - C
Zinc - Zn
Iron - Fe 
Hydrogen - H
Copper - Cu

Acid + Metal > salt + hydrogen - speed of reaction is indicated by rate of bubbles of hydrogen given off

Metal + water > Metal Hydroxide + hydrogen
Metals that will react with water: potassium, sodium, lithium and calcium
Metals that won’t react with water: zinc, iron and copper

Question 6

Separating metals from metal oxides

Answer 6

Metals often have to be separated from their oxides
Common metals react with oxygen to form oxides

Formation of metal ore :
Oxidation = gain of oxygen
E.g magnesium is oxidises to make magnesium oxide

Extraction of metal:
Reduction = loss of oxygen
E.g copper oxide is reduced to copper

Some metals can be extracted by reduction with carbon E.g
Iron oxide + Carbon > iron + carbon dioxide

The position of the metal in the reactivity series determines wether it can be extracted by reduction with carbon
Metals higher than Carbon have to be extracted via electrolysis
Metals below carbon can be extracted by reduction using carbon

Question 7

Redox Reactions

Answer 7

OILRIG - Oxidation is loss - Reduction is gain

A loss of electrons = oxidation
A gain of electrons = reduction

Displacement reactions are Redox reactions
A more reactive metal will displace a less reactive metal from its compound

Ionic equations - only the particles that react and products they form are shown

Question 8

Electrolysis

Answer 8

Splitting up with electricity

1) an electric current is passed through an electrolyte ( a molten or dissolved ionic compound). The ions move towards the electrodes where they react and the compound decomposes
2) the positive ions in the electrolyte will move towards the cathode ( negative electrode ) and gain electrons ( they are reduced
3) the negative ions in the electrolyte will move towards the anode ( positive electrode ) and lose electrons ( they are oxidised)
4) this creates a flow of charge through the electrolyte as ions travel to the electrodes
5) As ions gain or lose electrons, they form the uncharged element and are discharged from the electrolyte

Solids must be Molten before electrolysis can take place as they are in fixed positions.

Metals can be extracted from their ores using electrolysis

Question 9

Electrolysis of aqueous solutions

Answer 9

It may be easier to discharge ions from water than solute

Half equations - make sure electrons balance