Topic 2 Electrolysis

In a nutshell!

Indicators: the pH scale

Guess the pH

Your stomach contains hydrochloric acid. This aids digestion and allows the stomach to kill any harmful bacteria. The acid has a pH of ~2.

If to much acidic foods are eaten it can result in indigestion, heartburn and acid reflux. This causes pain in the stomach and in the oesophagus.

Alkali’s and Bases

The Facts!

Metal oxides

Ammonia An alkali is a base that dissolves in water.

OH Na

Sodium hydroxide

Cl H

Hydrochloric acid

The sodium replaces the hydrogen from HCl

Cl Na

Sodium chloride

H2O

Water

acid + base salt + water

Neutralisation

Particle model of Neutralisation

Hydrochloric acid

Sodium hydroxide

Sodium chloride

Water

•The technique of titration is used to find out accurately how much of a chemical substance is

dissolved in a given volume of a solution.

•The technique uses a particular set of apparatus with which volumes of solutions can be measured to an

accuracy of greater than 0.1cm 3

What is Titration?

Titration demonstration B Use the funnel to help you to pour

some hydrochloric acid into the burette. Put the

beaker under the burette and open the tap to let

some acid into the tip of the burette. Close the tap

again and pour the acid in the beaker down the sink. C Fill the burette up to just below the zero mark with hydrochloric acid – this does not have to

be exact. D Use the measuring cylinder to pour

exactly 10 cm3 sodium hydroxide solution into to

the conical flask. Add a few drops of indicator and

swirl it so that the liquids mix. Stand the flask under the burette. E Read the volume of acid in the burette and write it down. F Add the acid from the burette to the flask a little at a time – about 1 cm3 each time is

about right. Swirl the flask gently each time you

add some acid. Stop when the indicator turns red and stays red. G Record the volume of acid in the

burette. Work out how much acid has been added and write it down. H Wash out the conical flask with water

and repeat steps D to G. This time, add acid about

1 cm3 at a time until you have nearly added the

same amount as before. Then add acid in much smaller amounts – drop by drop.

Copper (II) oxide is a _________, which reacts with sulfuric acid to make a ________ and __________. A soluble base is called an ________. This is called a ______________ reaction.

Applying ideas…….

The Rules

Hydrochloric Acid → Chloride

Sulfuric Acid → Sulfate

Nitric Acid → Nitrate

Must remember……. METAL + ACID → SALT + HYDROGEN METAL OXIDE + ACID → SALT + WATER METAL CARBONATE + ACID → SALT + WATER + CARBON DIOXIDE

Zinc + hydrochloric + acid

• Use the general equation to finish of the word equations below metal + acid a salt + hydrogen

iron + nitric acid +

calcium + sulphuric

acid

+

iron nitrate

calcium sulphate

zinc chloride

hydrogen

hydrogen

hydrogen

magnesium + sulphuric

acid

+ magnesium sulphate

hydrogen

Activity

Metal oxides and acids

Metal oxide + acid a salt + water

• Most metal oxides are not soluble. • i.e. They are bases but not alkalis.

This means they are often slower and may need heating to make them react

Acid

Oxide

water + copper sulphate

sulphuric acid

+ Copper oxide

Reactions of metal oxides with acid A metal oxide is a compound containing a metal and oxide. They are sometimes called BASES. For example:

Mg O Na

Na O

O

Al

Al O

O

Magnesium oxide Sodium oxide Aluminium oxide

METAL OXIDE + ACID SALT + WATER

Copy and complete the following reactions:

1) Magnesium oxide + hydrochloric acid

2) Calcium oxide + hydrochloric acid

3) Sodium oxide + sulphuric acid

Mg O

H Cl

Mg Cl Cl

H H O H Cl

Sulfuric acid + copper carbonate

When your base is a carbonate, carbon dioxide

Gas is also made!!!

Base + acid a salt + water

1. Hydrochloric acid + sodium hydroxide

2. Hydrochloric acid + calcium carbonate

3. Sulfuric acid + copper oxide

4. Sulfuric acid + iron carbonate

5. Nitric acid + ammonia

6. Phosphoric acid + iron oxide

Task - write out and complete these half equations

C

1. Sodium chloride + water

2. Calcium chloride + water + carbon dioxide

3. Copper sulfate + water

4. Iron sulfate + water + carbon dioxide

5. Ammonium nitrate + water

6. Iron phosphate + water

Must remember……. METAL + ACID → SALT + HYDROGEN METAL OXIDE + SALT → SALT + WATER METAL CARBONATE → SALT + WATER + CARBON DIOXIDE

Complete the worksheet Neutralisation reactions

Extracting metals from ores Potassium

Sodium

Calcium

Magnesium

Aluminium

Carbon

Zinc

Iron

Tin

Lead

Copper

Silver

Gold

Platinum

Metals ABOVE CARBON, because of their high reactivity, are extracted by ELECTROLYSIS and this needs a lot of energy

Metals BELOW CARBON are extracted by heating them with carbon in a BLAST FURNACE

These LOW REACTIVITY metals blatantly won’t need to be extracted because they are SO unreactive you’ll find them on their own, not in a metal oxide

Fill in the gaps

• A solid compound cannot conduct _______. This is because the particle are not free to move. When we dissolve or melt the compound the _________ are free to move so can conduct ________ .

Copper sulphate

Copper ion

Sulphate on

Copper Sulphate solid

Electrolysis of copper sulphate

Negative electrode cathode

Positive electrode anode

electrolyte

+

Cation, copper

-

Anion, oxygen

+ -

Electrolysis Electrolysis is used to extract a HIGHLY REACTIVE metal.

= sulphate ion

= copper ion

When we electrolysed copper sulphate the

negative sulphate ions moved to the positive

electrode and the positive copper ions moved to the

negative electrode – OPPOSITES ATTRACT!!!

Electrolysis is splitting up substances using

electricity

(Lysis is latin for splitting)

What substances can we split up?

• We can separate compounds like: –Zinc chloride

–Lead bromide

–Aluminium oxide

–Nickel sulphate

They must be liquids

Write a sentence to explain what electrolysis is and write what will be made during electrolysis of these chemicals.

•Positive •Anode •Negative •Is •Cathode

Don’t get stressed in the exam: Remember PANIC

What happens at the electrodes

• At the positive electrode the negative anions are attracted to it. They give electrons to the electrode to lose their charge

+

-

e-

What happens at the electrodes

• At the negative electrode the positve cations are attracted to it. They take electrons from the electrode to lose their charge

-

+

e-

Electrolysis of HCl

+ -

- +

CATHODE

ANODE

H+

Cl-

H+

Cl-

H+

Cl- H+

Cl-

What is Test For Hydrogen?

Hydrochloric Acid

Magnesium Ribbon

POP!

What is the test for Chlorine?

Damp blue litmus paper Turns Then

Pink

White

Hydrogen chloride and hydrochloric acid have the formula HCl

Hydrogen chloride → hydrogen + chloride

HCl → H + Cl

HCl → H2 + Cl2

2HCl → H2 + Cl2

• Chlorine gas was used as a weapon in WW1.

• The gas dissolves the moisture in the eyes and lungs.

• The cholrine then forms an acid.

• The victim usually died very painfully within a few days.

Chlorine gas in the First World War

Can we use chlorine for anything?

PVC

Paper making

Electrolysis of brine

• When NaCl dissolves in water, it’s ions become free to move. So the solution can be electrolysed.

• In water some of the molecules of water will naturally split apart.

Watch the demonstration of the electrolysis of brine

Product Test for substance

Hydrogen Blows out a lighted splint

and makes a squeaky

pop.

Chlorine Damp litmus paper goes

from blue to red then

bleaches white

Sodium hydroxide Goes dark blue or purple

when universal indicator

is added.

Brine, Sodium Chloride in water

sodium ion

chlorine on

Sodium chloride solid

Sodium chloride

sodium ion

chloride on

Sodium chloride solid

Electrolysis of NaCl solution

Electrolysis of salt 1

Cl2 2Cl- - + 2e- Chlorine gas is formed

Negative ions

• Salt consists of sodium ions (Na+) and chloride ions (Cl-).

• Chloride ions go to the anode where they lose an electron

• The neutral chlorine atoms produced join up into pairs

Electrolysis of salt 3

• Na+ ions move to the cathode but do not accept electrons.

• It is the hydrogen ions that gain electrons

• As a result hydrogen gas is formed at the cathode.

2H+ + 2e- H2

How does the sodium hydroxide form? Sodium chloride solution has four types of ions:

What is the overall equation for the electrolysis of a sodium chloride solution?

The Cl- ions form chlorine at the positive electrode and the H+ ions form hydrogen at the negative electrode. So, what’s left?

Na+ and OH- ions are left behind and so a solution of sodium hydroxide (NaOH) is formed.

2NaCl (aq) + 2H2O (l) H2 (g) + Cl2 (g) + 2NaOH (aq)

Na+ and Cl- ions from the sodium chloride H+ and OH- ions from the water.

Draw what you did

Power pack

Beaker

Test-tube

Sodium + water→ sodium + hydrogen + Chlorine

Chloride hydroxide

NaCl + H2O → NaOH + H2 + Cl2

2NaCl + 2H2O → 2NaOH + H2 + Cl2

What would we get if we electrolyse water? How much of each gas would we expect to get? H2O

Electrolysis of dilute sulfuric acid

Uses of hydrogen

Rocket fuel

Uses of oxygen

Welding equipment

Compound Anode Cathode

Sodium Bromide

Potassium Iodide

Calcium Fluoride

Magnesium Oxide

Lithium Chloride

O.I.L.R.I.G.

•Oxidation is loss of electrons

•Reduction is gain of electrons

Half Equations

• The half equation shows what happens at each electrode e.g.

3Al3+ + 3e- 3Al • The aluminium ions collect 3 electrons at the

cathode to form aluminium atoms. 2Cl- Cl2 + 2 e-

• The chlorine ions drop off electrons at the anode to form chlorine atoms (which in turn react to form a chlorine molecule (covalent bonding)

Balance the following half equations.

1. Cl2 + e- → Cl-

2. Al3+ + e- → Al

3. H+ + e- → H2

4. Pb2+ + e- → 2Pb

5. O2- - 4e- → O2

Watch the animation on page 119 of the active book

The Reactivity Series

The Reactivity Series lists metals in order of reactivity:

Make a pneumonic to help you remember this

Potassium

Sodium

Calcium

Magnesium

Aluminium

Carbon

Zinc

Iron

Lead

Copper

Silver

Gold

The premier league of reactivity •Think about it like a football league. If the team near the bottom played the team near the top, who would

you expect to win?? It’s the same for the reactivity series, if an element

near the top of the reactivity series is mixed with an ore from the bottom, the element displaces the less

reactive metal from it’s ore.

V

Explain why we need different methods of extraction depending on the reactivity of the metal Objective:

Displacement reactions

Mg

Magnesium

SO4 Cu

Copper sulphate

The magnesium DISPLACES the copper from copper sulphate

SO4 Mg

Magnesium sulphate

Cu

Copper

A displacement reaction is one where a MORE REACTIVE metal will DISPLACE a LESS REACTIVE metal from a compound.

Magnesium + copper sulphate magnesium sulphate + copper

Watch the dating demo

Some example reactions… Reaction Prediction Observations

Zinc + copper sulphate

Zinc + lead nitrate

Copper + lead nitrate

Copper + silver nitrate

Extension work – write down the equations for these reactions

Zinc

sulphate

Copper

sulphate

Iron

sulphate

Magnisium

sulphate

Zinc

X Copper

X Iron

X magnesium

X

4. Aisha placed small samples off four different metals on a spotting tile. She added drops of copper sulphate solution to each metal.

Aisha repeated the experiment with fresh samples of the four metals and solutions of different salts. She recorded some of her results in a table. shows that a reaction took place X shows that no reaction took place.

copper iron magnesium zinc

spotting tile

coppersulphate

ironsulphate

magnesiumsulphate

zincsulphate

solutionsmetals

copper iron magnesium zinc

×

×

×

× ×

×

×

×

(a)The four metals have different reactivities. (i) Use the information in the table to put the four metals in a reactivity series. 1 mark (ii) Use the reactivity series to complete the table by writing in √ or X in the three empty boxes. 2 marks (b) Copper reacts with silver nitrate solution. (i)Complete the word equation for the reaction: Copper + silver nitrate → ...........................+........................... 2 marks (ii) Platinum does not react with silver nitrate.

Put the metals platinum, copper and silver in the correct order according to their reactivity. (iii) 1 mark (c) In many houses the hot water pipes are made from copper and the boiler is made from iron. Which of these metals will corrode first? Explain your answer. 1 mark Maximum 7 marks

Rocks and Minerals

Objective: Recall that most metals are extracted as ores from the earth’s crust

•Rocks in the earth’s crust are very rarely pure substances •If the minerals in rocks contain metals then they are called ores. •Unreactive metals are found in their pure state and not as an ore.

•Can you think of any unreactive metals? •Metals vary in cost depending on the availability of the ore, how much metal is contained in the ore and the method of extraction. •You are going to try one method of extraction today.

bauxite(Aluminium)

haematite(iron)

cinnabar(mercury)

Galena(lead)

malachite(copper)

Match the metal ore to the picture

Starter Activity

bauxite(Aluminium)

haematite(iron) cinnabar(mercury)

Galena(lead)

malachite(copper)

Match the metal ore to the picture

Answers

Which method of extraction?

Objective: Explain why we need different methods of extraction depending on the reactivity of the metal

•There are two main methods for the extraction of metals from their ores: •Chemical reduction (like you’ve carried out in class) •Electroysis •The method used for extraction depends on the reactivity of the metal you are trying to extract. •Scientists use something called the reactivity series to help them decide which method to use, the higher the metal, the more reactive it is.

Extracting metals from ores Potassium

Sodium

Calcium

Magnesium

Aluminium

Carbon

Zinc

Iron

Tin

Lead

Copper

Silver

Gold

Platinum

Metals ABOVE CARBON, because of their high reactivity, are extracted by ELECTROLYSIS and this needs a lot of energy

Metals BELOW CARBON are extracted by heating them with carbon in a BLAST FURNACE

These LOW REACTIVITY metals blatantly won’t need to be extracted because they are SO unreactive you’ll find them on their own, not in a metal oxide

Oxidation

A reaction where the products contain more oxygen than the original fuel. Adding more oxygen to a substance is called ‘oxidation’.

Gloss paint Iron nails

Burning

Oxidation Fuels like carbon when burnt are oxidised.

C CO2

S SO2

It has had oxygen added to it.

Iron has been oxidised to make iron oxide (rust).

Reduction

Removing oxygen from a chemical is called reduction. It reduces the amount of oxygen in the substance.

Reduction Reduction can occur when burning iron oxide.

3C + 2Fe2O3 4Fe + 3Co2

The iron oxide gives its oxygen to the carbon.

What can we say has happened to the Carbon?

Redox reactions

Although some fuels gain oxygen some will also lose oxygen. Thus oxidation and reduction are happening at the same time. These are known as REDOX reactions.

Write a word equation for methane (CH4) reacting with oxygen. State what where redox

is occurring.

NaOCl NaCl CH3CH2OH CH3COOH

Oxidation or Reduction? How do we know?

What are the properties of metals

What are the

properties of non- metals

Comparing the electrical conductivity of metals:

• Are all metals good conductors?

• Are some better than others?

• How could we investigate this?

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What are the advantages of recycling metals?

• Make metals last longer • Less energy needed to

recycle compared to extracting ore

• Reduces the need to mine ores – less damage to the environment

• Less pollution produced • Less waste materials put

into landfill

Monday, 16 March 2015

What are the disadvantages of recycling metals?

• Costs and energy

involved in:

– Collecting materials

– Sorting materials

– Transporting

materials

• Can be more expensive to recycle

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The graph…

Monday, 16 March 2015

An alloy is a mixture of a metal with at least one other element.

The final alloy may have very different properties to the original metal.

What is an alloy?

Steel is a common example of an alloy. It contains iron mixed with carbon and other elements. Adding other elements to a metal changes its structure and so changes its properties.

What types of alloys are there?

brass: an alloy of copper and zinc.

solder: an alloy of tin and lead

amalgam: an alloy of mercury and

silver or tin.

Other well-known alloys include:

Steel is an alloy of iron and other elements, including carbon, nickel and chromium.

Steel is stronger than pure iron and can be used for everything from sauce pans… …to suspension bridges!

What is steel?

When other elements are added to iron, their atoms distort the regular structure of the iron atoms.

The atoms in pure iron are arranged in densely-packed layers. These layers can slide over each other. This makes pure iron a very soft material.

Why is steel stronger than iron?

It is more difficult for the layers of iron atoms in steel to slide over each other and so this alloy is stronger than pure iron.