Name: _________________________ 

Teacher: ______________ 

EXAM DATE:  

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Oxidation and reduction

In oxidation reactions, a substance often gains oxygen.

In reduction reactions a substance often loses oxygen.

Oxidation and reduction always occur together. Usually, when one substance is oxidised, another is reduced at the same time.

Metals react with oxygen to form metal oxides. For example, when magnesium is burned in air, it reacts with oxygen to form magnesium oxide. The magnesium gains oxygen in the reaction, so it is oxidised.

Magnesium + Oxygen Magnesium oxide

2Mg + O2 2MgO

Metal oxides can be reduced by removing oxygen. For example, when lead (IV) oxide is heated with carbon:

- The lead (IV) oxide loses oxygen, so it is reduced. - The carbon gains oxygen, so it is oxidised.

Lead(IV) oxide + Carbon Lead + Carbon dioxide

PbO2 + C Pb + CO2

The reactivity series

You must know the reactivity 

series, and which elements are 

more or less reactive than others. 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Displacement reactions

In a displacement reaction, a more reactive metal will displace a less reactive metal from a solution of its salt. Magnesium is more reactive than copper, so magnesium will displace copper from a solution of copper sulfate.

Magnesium + Copper Sulfate Copper + Magnesium sulfate

Mg(s) + CuSO4(aq) Cu(s) + MgSO4(aq)

Less reactive metals will not displace more reactive metal salts.

The reactivity series (continued)

When metals react, their atoms lose electrons to form positive ions.

Some metals can lose their electrons more easily than others.

The more easily a metal loses its electrons, the more reactive it is. For example, potassium loses an electron easier than sodium loses an electron, therefore potassium is more reactive.

The reaction of metals with acid and water can be used to place them in order of reactivity. This is called the reactivity series.

Metals react with acids to produce metal salts and hydrogen.

Lithium, sodium and potassium are very reactive – they react vigorously with water to produce a metal hydroxide solution and hydrogen. These metals are placed at the top of the reactivity series. Lithium, sodium and potassium would react so vigorously with dilute acid that it would not be safe to carry out these reactions.

Calcium, magnesium, zinc and iron are fairly reactive metals – they react quickly with acid and slowly with water.

Very unreactive metals, like copper and gold, do not react with acids or water and are placed at the bottom of the reactivity series.

Carbon and hydrogen are included in the reactivity series are also added for comparison.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

   

Extraction of metals

The method of extraction of a metal depends on how reactive it is. Unreactive metals (e.g. gold) exist as elements at the Earth’s surface. These are called native metals. However, most metals are found as metal oxides, or as compounds which can be easily changed into metal oxides.

Metals that are less reactive than carbon (e.g. iron and lead) can be extracted from their oxides by heating with carbon:

Iron oxide + Carbon Iron + carbon dioxide

2Fe2O3 + 3C 4Fe + 3CO2

The iron oxide loses oxygen, so it is reduced.

The carbon gains oxygen, so it is oxidised.

Losing or gaining electrons.

This is for higher tier ONLY

Not all reduction and oxidation reactions involve oxygen.

Because of this, scientists use the following acronym. Remember this!

Oxidation Is Loss (of electrons) Reduction Is Gaining (of electrons)

The balanced symbol equation for the reaction between magnesium and oxygen can be split into two ionic equations:

2Mg 2Mg2+ + 4e-

O2 + 4e- 2O2-

This shows that magnesium loses electrons, and oxygen gains electrons.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

   

Quick test!

1. Calcium + oxygen calcium oxide Which substance is oxidised in this reaction. ____________________________________________________ ____________________________________________________

2. Complete the word equation: Iron + Copper sulfate _____________ + _____________

3. By what method should lead be extracted from lead oxide? ____________________________________________________ ____________________________________________________

This is for higher tier ONLY

4. Calcium reacts with hydrochloric acid to form calcium chloride and hydrogen. The equation for this reaction can be split into two ionic equations. Ca Ca2+ + 2e- 2H+ + 2e- H2

a) Which species is oxidised? Explain your answer. ________________________________________________________________________________________________________________________________________________

b) Which species is reduced? Explain your answer. ________________________________________________________________________________________________________________________________________________

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

The pH scale

When substances dissolve in water, they dissociate into their individual ions:

- Hydroxide ions, OH-(aq) make solutions alkaline. - Hydrogen ions, H+(aq) make solutions acidic.

The pH scale is a measure of the acidity or alkalinity of any aqueous solution:

- A solution with a pH of 7 is neutral. - Aqueous solutions with a pH of less than 7 are

acidic. - The closer the pH is to 0, the stronger the acid - Aqueous solutions with a pH of more than 7 are

alkaline. - The closer the pH is to 14, the stronger the alkali.

The pH of a solution can be measured using a pH probe or universal indicator.

Indicators are dyes that change colour depending on whether they are in acidic or alkaline solutions:

- Litmus changes colour from red to blue, or vice versa. - Universal indicator is a mixture of dyes that shows a range of colours to

indicate how acidic or alkaline a solution is. This colour is then compared to the pH scale.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

   

Neutralisation of acids

Bases are substances which can neutralise acids. If the base can dissolve, we call it an ‘alkali’.

Acid + Metal hydroxide (alkali) Salt + Water

Acids contain hydrogen ions, H+(aq) Alkalis contain hydroxide ions (OH-) When an acid reacts with an alkali, the H+ ions and OH- ions react together to produce water, H2O, which has a pH of 7.

H+(aq) + OH-(aq) H2O(l)

This type of reaction is called a neutralisation reaction. - Acid is neutralised by the alkali - The solution that remains has a pH of 7, showing it is neutral.

Acids can also be neutralised by metal oxides and metal carbonates. These are bases, as they cannot dissolve (they are insoluble).

Acid + Metal oxide Salt + Water Acid + Metal carbonate Salt + Water + Carbon dioxide

A salt is produced when the hydrogen in the acid is replaced by a metal ion.

HCl + NaOH NaCl + H2O

The name of the salt produced depends on the name of the acid used.

Acid Name of salt Example Hydrochloric acid … Chloride Sodium chloride (NaCl)Sulfuric acid … Sulfate Coper sulfate (CuSO4)Nitric acid … Nitrate Potassium nitrate (KNO3)

 

 

 

 

 

   

Required practical – preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate.

Sample method: 1. Add the metal oxide or carbonate to a warm

solution of acid until no more will react. 2. Filter the excess metal oxide or carbonate to

leave a solution of the salt. 3. Gently warm the salt solution so that the

water evaporates and crystals of salt are formed.

Hazards and risks: - Corrosive acid can cause

damage to eyes, so eye protection must be worn.

- Hot equipment can cause burns so care must be taken when the salt solution is warmed.

Strong and weak acids

This is for higher tier ONLY

Strong acids are completely ionised (split up into their ions) in water.

Hydrochloric acid is a strong acid:

HCl(g) + aq H+(aq) + Cl-(aq)

Ethanoic acid is a weak acid, so is not completely ionised.

CH2COOH(l) + aq CH3COO-(aq) + H+(aq)

The pH of a solution is a measure of the concentration of the H+ ions. A pH decrease of one unit indicates that the concentration of H+ ions has increased by a factor of 10. A stronger acid will have more H+ ions, and will therefore have a lower pH.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Quick test!

1. What is the pH of a neutral solution? ____________________________________________________

2. Complete the general equation below: Acid + Alkali ________ + ________

3. Describe a quick method for the preparation of a sample of pure and dry sodium chloride salt. ____________________________________________________________________________________________________________________________________________________________

This is for higher tier ONLY

4. What is a strong acid? ________________________________________________________________________________________________________

5. The pH of a solution changes from 6 to 4. What happens to the concentration of hydrogen ions? ____________________________________________________________________________________________________________________________________________________________

Electrolysis

Electrolysis is the use of an electrical current to break down compounds containing ions into their constituent elements. The substance being broken down is called the electrolyte. The electrodes are made from solids that conduct electricity. During electrolysis: - Negatively charged ions move to the anode

(positive electrode). - Positively charged ions move to the cathode

(negative electrode).

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Electrolysis (continued)

Electrolysis can be used to separate ionic compounds into elements. For example, lead bromide can be split into lead and bromine:

- The lead bromide is heated until it melts. - The positively charged lead ions move to the negative electrode

(cathode). - Here they gain electrons to form lead atoms – pure lead is produced at

this electrode. - The negatively charged bromide ions move to the positive electrode

(anode). - Here they lose electrons to form bromine atoms, which join together to

form bromine molecules – bromine is released at this electrode.

At the cathode: Pb2+ + 2e- Pb At the anode: 2Br- Br2 + 2e-

Ionic substances can only conduct electricity when they are molten (melted) or dissolved in water.

Oxidation and reduction

This is for higher tier ONLY

Reduction occurs when positively charged ions gain electrons at the negative electrode (anode). Oxidation occurs when negatively charged ions lose electrons at the negative electrode (cathode). In a redox (reduction and oxidation) reaction which takes place at the electrodes during electrolysis, the reaction can be represented during half equations. For example, in the electrolysis of molten copper chloride: - Copper is deposited at the negative electrode

Cu2+ + 2e- Cu - Chlorine gas is given off at the positive electrode:

2Cl- Cl2 + 2e-

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Extraction of metals

Metals that are more reactive than carbon can extract from their ores using electrolysis.

Electrolysis requires lots of heat and electrical energy, making it an expensive process.

Aluminium is obtained by the electrolysis of aluminium oxide that has been mixed with cryolite (a compound of aluminium).

Cryolite lowers the melting point of aluminium oxide, meaning less energy is needed (cheaper energy costs).

Aluminium forms at the negative electrode.

Oxygen forms at the positive and reacts with carbon to form carbon dioxide. This wears away the positive electrode, which is replaced regularly.

At the cathode: Al3+ + 3e- Al

At the anode: 2O2- O2 + 4e-

Electrolysis of aqueous solutions

When ionic compounds are dissolved in water to form aqueous solutions, it is slightly harder to predict the products of electrolysis. The water molecules break down to form hydroxide ions, OH-, and hydrogen ions, H+.

At the negative electrode: - Hydrogen is produced if the metal is more reactive than hydrogen. - The metal is produced if the metal is less reactive than hydrogen.

At the positive electrode: - Oxygen is produced unless the solution contains halide (group 7) ions. - If halide ions are present, then the halogen is produced.

In the electrolysis of sodium chloride: - Hydrogen is released at the negative electrode. - Chlorine gas is released at the positive electrode.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Required practical: Investigate what happens when aqueous solutions are electrolysed using inert electrodes.

Sample method: 1. Set up the equipment

as shown in the diagram under section “Electrolysis”

2. Pass an electric current through the aqueous solution.

3. Observe the products formed at each inert electrode.

Hazards and risks- A low voltage must be

used to prevent an electric shock.

- The room must be well ventilated, and the experiment must only be carried out for a short period of time, to prevent exposure to dangerous amounts of chlorine gas.

Quick test!

1. Give the scientific names of the positive electrode and negative electrodes. ________________________________________________________________________________________________________

2. Suggest which charge of ions move to the positive and negative electrode. ________________________________________________________________________________________________________

3. Predict the products of the electrolysis of aqueous sodium bromide. ________________________________________________________________________________________________________

This is for higher tier ONLY

4. Write down the half equations for the reactions that take place at each electrode in the electrolysis of aqueous sodium bromide. ____________________________________________________________________________________________________________________________________________________________

 

   

Practice exam questions

This is FOR ALL ABILITIES

1. Calcium can be burned to produce calcium oxide: 2Ca + O2 2CaO

Write the word equation for this reaction (2 marks) _____________________________________________________

2. The list shows the reactivity of metals. Which two statements about this series are correct? Tick two boxes (2 marks) Copper is more reactive than potassium Sodium is more reactive than magnesium Magnesium is more reactive than aluminium Aluminium is less reactive than iron

3. Complete the displacement reactions below (6marks): Aluminium + Iron oxide __________ + __________ Magnesium + ________ dioxide ______________ + Carbon Copper oxide + Gold __________ + __________

4. Identify the type of reaction which involves the removal of oxygen (1

mark) _____________________________________________________

5. Use the reactivity series to describe which elements can be extracted using carbon, but not by using hydrogen (2 marks) _____________________________________________________ _____________________________________________________

6. Explain how carbon can be used to extract some metals from their metal oxides (3 marks) _____________________________________________________ _____________________________________________________ _____________________________________________________

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Practice exam questions

1. State the pH value of a solution with a neutral pH (1 mark) _____________________________________________________

2. State the range of pH of a solution with an acidic pH (1 mark) _____________________________________________________

3. State the range of pH of a solution with an alkali pH (1 mark) _____________________________________________________

4. Describe how to test to see whether a solution is acidic, alkali or neutral (3 marks) _____________________________________________________ _____________________________________________________ _____________________________________________________

5. Complete the equation to show the names of the salts produced (3 marks) Hydrochloric acid + Sodium hydroxide Water + ________________Sulfuric acid + Lithium oxide Water + _______________________Nitric acid + Potassium hydroxide Water + ___________________

6. Give the ionic equation when acids react with alkalis (3 marks) _____________________________________________________

7. Describe a method on how to safely obtain a pure and dry sample of salt (4 marks). _____________________________________________________ _____________________________________________________ _____________________________________________________ _____________________________________________________ __________________________________________________________________________________________________________ _____________________________________________________

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Practice exam questions

1. Give the scientific names of the positive electrode and negative electrodes (2 marks). __________________________________________________________________________________________________________

2. Copper chloride is an ionic compound that can be separated by electrolysis. Name the element formed during the electrolysis of molten copper chloride:

i. At the positive electrode (1 mark) ______________ ii. At the negative electrode (1 mark) ______________

3. A student carried out an experiment to find out what happens when an aqueous solution of sodium chloride is electrolysed.

i. Identify the two positive ions present in an aqueous solution of sodium chloride. (2 marks) _________________________________________________

ii. Name the substance produced at the negative electrode. You must give a reason for your answer. (2 marks) _________________________________________________

iii. Name the substance produced at the positive electrode. You must give a reason for your answer. (2 marks) _________________________________________________

4. During the electrolysis of lead chloride, lead and chlorine are produced.At the anode: 2Cl- Cl2 + 2e- At the cathode: Pb2+ + 2e- Pb Explain why the lead chloride has to be molten before electrolysis can occur (2 marks) _______________________________________________________________________________________________________________________________________________________________

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

   

 

 

 

 

 

Practice exam questions

This is FOR HIGHER TIER ONLY

1. The balanced symbol equation for the reaction between oxygen and calcium can be broken down into two ionic equations:

2Ca 2Ca2+ + 4e- O2 + 4e- 2O2-

In terms of oxidation and reduction, explain what happens to the calcium and the oxygen in the reaction (4 marks) __________________________________________________________________________________________________________ __________________________________________________________________________________________________________

2. Hydrochloric acid is a strong acid while carbonic acid is a weak acid. Explain the difference between a strong and a weak acid (2 marks). __________________________________________________________________________________________________________ _____________________________________________________

3. A 0.01mol/dm3 solution of HCl has a pH of 2. Predict the pH of a 0.001mol/dm3 solution of hydrochloric acid (1 mark) _____________________________________________________

4. When aqueous potassium nitrate solution is electrolysed, neither potassium nor nitrogen is discharged. Explain why and state what is actually produced instead (4 marks) __________________________________________________________________________________________________________ _____________________________________________________

5. Write two half equations for the reaction that occurs when water is electrolysed. At the cathode: _______________________ At the anode: _________________________