Unit 3

Acids, Bases and Metals

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• Acids and bases• Salt preparation• MetalsClick here to End.

Acids and Bases

pH

• pH is a scale of acidity.• It can be measured using:• pH paper• Universal Indicator solution • A pH meter.

pH

•  pH is a continuous scale of acidity which runs from below 0 to above 14.

• Acids have a pH of less than 7 • Alkalis have a pH of more than

7. • Pure water, and other neutral

solutions have a pH equal to 7.

Oxides

• Non-metal oxides which dissolve in water, give acid solutions.

• Metal oxides and hydroxides, which dissolve in water, give alkaline solutions.

• Ammonia dissolves in water to produce an alkali.

• Acids and alkalis are in common use at home or in the laboratory.

Ions

• Acids and alkalis both contain ions.

• Acids contain the H+(aq) ion• Alkalis contain the OH-(aq) ion• In water the concentration of

ions is very low.

H+ and OH- ions

• Water, and other neutral solutions, contain equal numbers of H+ and OH- ions.

• Acidic solutions contain more H+ ions than OH- ions .

• Alkalis contain more OH- ions than H+ ions.

Dilution

•  When an acid is diluted its acidity decreases and its pH increases.

• When an alkali is diluted its alkalinity decreases and its pH decreases.

•  When an acid (or alkali) is diluted then the number of H+ (or OH- ) ions per cm3 of solution decrease and so the acidity (or alkalinity) decrease.

Equilibrium

• There is an equilibrium in water:

• H2O(l) H+(aq) + OH-(aq)

• This means that the concentrations of reactants and products are always the same (but not equal).

Concentration

• The concentration of a solution is measured in moles per litre (mol l-1)

• A 1 mol l-1 solution means 1 mole is divided in each litre of solution.

• To connect the concentration of a solution to the number of moles and concentration use the triangle opposite.

• c = concentration (m/l)

• n = number of moles

• v = volume (l)

n

vc

Strong and Weak Acids

• A strong acid is one which completely dissociates in water:HCl(aq) H+(aq) + Cl-(aq)

• A weak acid is one which partially dissociates in water:

CH3CO2H(aq)H+(aq) + CH3CO2-

(aq)

Strong Acids

• Hydrochloric acid, nitric acid and sulphuric acid are strong acids.

HCl(aq) H+(aq) + Cl-(aq)HNO3(aq) H+(aq) + NO3

-(aq)

H2SO4(aq) 2H+(aq) + SO42-

(aq)

Weak Acids

• Ethanoic acid, is a weak acid.CH3CO2H(aq)H+(aq) + CH3CO2

-

(aq)

Comparing Weak and Strong Acids

Test 100 ml 0.1 mol/l HCl

100 ml 0.1 mol/l CH3CO2H

pH 1 3

Conductivity

Very high Low

Rate of reaction

Fast Slow

Strong and Weak Bases

• A strong base is one which completely dissociates in solution:NaOH(aq) Na+(aq) + OH-

(aq)• A weak base is one which

partially dissociates in solution:NH4OH(aq)NH4

+(aq) + OH-(aq)

Strong Bases

• Solutions of metal hydroxides are strong basesNaOH(aq) Na+(aq) + OH-(aq)

KOH(aq) K+(aq) + OH-(aq)

Weak Bases

• A solution of ammonia is a weak base.

NH3(g) + H2O(l) NH4OH(aq)

NH4OH(aq) NH4+(aq) + OH-

(aq)

Comparing Weak and Strong Bases

Test 100 ml 0.1 mol/l NaOH

100 ml 0.1 mol/l NH4OH

pH 13 11

Conductivity

Very high Low

Rate of reaction

Fast Slow

Acids and Bases

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Salt Preparation

Neutralisation

• Neutralisation is the reaction of an acid with a base.

• Metal oxides, hydroxides and carbonates are all examples of bases.

• Bases which dissolve in water are called alkalis.

• During a neutralisation reaction then the pH of the acid involved moves up nearer to 7.

• During a neutralisation reaction then the pH of the alkali involved moves down nearer to 7.

• In the reaction of an acid and an alkali the hydrogen ions and hydroxide ions form water.H+ + OH- H2O

HCl + NaOH NaCl + H2O

• In the reaction of an acid and a metal oxide the hydrogen ions and oxide ions form water.2H+ + O2- H2O

H2SO4 + CuO CuSO4 + H2O

• In the reaction of an acid and a metal carbonate the hydrogen ions and carbonate ions form carbon dioxide and water.2H+ + CO3

2- CO2 + H2O

2HNO3 + CaCO3 Ca(NO3)2 + CO2 + H2O

• Examples of neutralisation involve adding lime to soil or water to reduce its acidity treating acid indigestion with magnesium hydroxide the reaction of H+ (aq) to form water.

Acids and metals

• Acids react with some metals to release hydrogen. The hydrogen ions in the the acid form hydrogen molecules.

• The test for hydrogen is that it burns with a “pop”.

Acid Rain

• Sulphur dioxide is produced by the burning of fossil fuels.

• Nitrogen dioxide is produced by the sparking of air in car engines.

• Both these gases dissolve in water in the atmosphere to produce acid rain.

• Acid rain has damaging effects on buildings made from carbonate rock, structures made of iron and steel, soils and plant and animal life.

Acids and carbonates

• An acid reacts with a metal carbonate to release carbon dioxide. Thus acid rain will dissolve rocks or buildings which contain carbonates.

• The hydrogen ions from the acid react with the carbonate ions, to form carbon dioxide and water. 2H+ + CO3

2- H2O + CO2

Remember Moles?

• To connect gram formula mass, mass in grams and number of moles use the triangle opposite

• gfm = mass of 1 mole

• n = number of moles• m = mass of

substance

m

gfmn

Remember solutions?

• To connect volume, concentration and the number of moles in a solution use the triangle opposite.

• c = concentration (m/l)

• n = number of moles

• v = volume (l)

n

vc

Working out about neutralisations

• Work out unknown concentrations and volumes from the results of volumetric titrations.

• You use the equation VH MH NH

=VOH MOH NOH

V = volume

M = molarityNH = number of H+ ions in acidNOH =number of OH- ions in alkaliH = acidOH = alkali

Salts

• A salt is the compound formed when the hydrogen ion of an acid is replaced by a metal ion (or an ammonium ion).

• Salts are formed by the reactions of acids with bases or metals.

Salts

Acid Formula

Salt Ion

hydrochloric

HCl chloride Cl-

sulphuric H2SO4 sulphate SO42-

nitric HNO3 nitrate NO3-

carbonic H2CO3 carbonate

CO32-

Making salts

• There are three main methods of making salts.

• Which method to use depends upon solubilities.

• Those solubilities can be found in the Data Booklet.

Titration

• Titration is experiment where alkali is measured out using a pipette.

• Indicator is added.• Acid is added from a burette,

until the indicator changes colour.

• The water can then be evaporated to get the salt.

Neutralisation and Evaporation.

• An easy way to prepare salts is to react an acid with an insoluble metal oxide or metal carbonate.

• The excess can be removed from the reaction mixture by filtration.

• The solution is now evaporated to separate the salt.

Precipitation

• Precipitation is the reaction in which two solutions react to form an insoluble salt.

• The salt can then be filtered out and dried.

Is salt soluble?

Yes

No

Precipitation

Is base soluble?

Yes

No

Titration

Neutralisation andEvaporation

How to make a salt.

Ionic equations

• Normally when we write equations we do so like this:

HNO3 + NaOH NaNO3 + H2O

• We can change to write the equations using ions.

H+ + NO3- + Na+ + OH-

Na+ + NO3- + H2O

• If we look closely we can see that some ions appear unchanged on both sides. We call these spectator ions.

H+ + NO3- + Na+ + OH-

Na+ + NO3- + H2O

• We can rewrite the equation, without the spectator ions.

• This shows the ions that participate in the reaction.

H+ + OH- H2O

Salt Preparation

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Metals

Cells

• Chemical changes can bring about the production of electrical energy.

• The current measures how many electrons flow in the cell each second.

• The voltage measures how hard those electrons are pushed by the chemicals used.

• A cell is made by connecting two different metals together with an electrolyte.

• An electrolyte is a material, which conducts electricity in solution (it contains ions). The electrolyte is needed to complete the circuit.

The Electrochemical Series

• We can measure the voltage produced by connecting different metals together to form a simple cell.

V

The Electrochemical Series

• The voltage between different pairs of metals varies.

• By listing the metals according to the voltage they produce we get the electrochemical series.

K

Na

Ca

Mg

Al

Zn

Fe

Sn

Pb

Cu

Hg

Ag

Au

The electrochemical series shows a “league table”.

The substances at the top are those best at pushing electrons.

Displacement

• Any metal, in an Electrochemical Series, will displace a metal lower down in the electrochemical series it from one of its compounds.

• Iron will displace copper from copper(II) sulphate solution.

• K• Na• Al• Zn• Fe• Sn• Pb• Cu

• Iron can displace copper from copper(II) sulphate solution.

• K• Na• Al• Zn• Fe• Sn• Pb• Cu

CuSO4Fe

Fe SO4Cu

• Iron cannot displace zinc from zinc(II) sulphate solution.

• K• Na• Al• Zn• Fe• Sn• Pb• Cu

ZnSO4Fe

ZnSO4Fe

• Most displacement reactions will give some visible signs.

• If zinc reacts with copper(II) sulphate solution:

• The silver colour of zinc will be replaced by the brown colour of copper.

• The blue colour of the solution will fade as Cu2+ changes to Cu.

Hydrogen

• By considering the metals with which acids will react it is possible to place hydrogen in the Electrochemical Series.

K

Na

Ca

Mg

Al

Zn

Fe

Sn

Pb

Cu

Hg

Ag

Au

K

Na

Ca

Mg

Al

Zn

Fe

Sn

Pb

Cu

Hg

Ag

Au

Metals which react with acid, releasing hydrogen

K

Na

Ca

Mg

Al

Zn

Fe

Sn

Pb

Cu

Hg

Ag

Au

Metals which react with acid, releasing hydrogen

Metals which donot react with acid.

K

Na

Ca

Mg

Al

Zn

Fe

Sn

Pb

Cu

Hg

Ag

Au

Metals which react with acid, releasing hydrogen

Metals which donot react with acid.

H

More about cells

• Electricity can be produced in a cell by connecting two different metals in solutions of their metal ions.

• The next slides show how copper and zinc half-cells can be made to make a cell.

Put a copper rodin a solution containingcopper ions.

Put a zinc rodin a solution containingzinc ions.

V

Connect the two metal rods through a voltmeter

There is no reading because the circuit is not complete.

V

Add an ion bridge (also called a salt bridge). This is to allow the movement of ions to complete the

circuit.

• Electricity can be produced in a cell when at least one of the half-cells does not involve metal atoms.

• Electrons flow through the meter from the substance higher in the electrochemical series to the one lower in the electrochemical series.

Oxidation

• Oxidation is the loss of electrons by a reactant in a chemical reaction.

• When a metal reacts to form a compound it is an example of oxidation.

Reduction

•  Reduction is the gain of electrons by a reactant in a chemical reaction.

• When a metal compound reacts to form a metal it is an example of reduction.

Oxidation and Reduction

•OIL RIGOxidation Is Loss of electrons Reduction Is Gain of electrons

• In a redox reaction oxidation and reduction go on together.

Redox

•Ion-electron equations can be written for oxidation and reduction reactions.

•These equations can be combined to produce redox equations.

• Magnesium can be oxidised.Mg Mg2+ + 2e oxidationHydrogen ions can be reduced.2H+ + 2e H2 reduction

Overall2H+ + Mg H2 + Mg2+ redox

• In the displacement of copper by zinc the reaction is:

• CuSO4 + Zn ZnSO4 + Cu

• This can be written as:Cu2+ + 2e Cu reduction Zn Zn2+ + 2e oxidation OverallCu2+ + Zn Cu + Zn2+ redox

Redox and Electrolysis

•Electrons are released at the negative electrode so reduction takes place there.

•Electrons are taken in at the positive electrode so oxidation takes place there.

Common reactions of metals

• Metals react with oxygen to form metal oxides.

• Metals react with water (either as liquid or steam) to form the metal hydroxide and hydrogen.

• Metals react with dilute acid to release hydrogen.

Reactions of Metals

• N.B. Not all metals react as shown on the previous slide.

• The ease with which these reactions take place is a measure of the reactivity of the metal.

• We can build up a Reactivity Series from the relative reactivity of the metals.

Recovering Metals

• Ores are naturally occurring compounds of a metal.

• Less active metals, such as silver and gold, do not react well and so occur uncombined in the earth's crust.

• They were the first metals discovered.

• The extraction of a metal from its ore is an example of reduction.

• Oxides of reactive metals are most difficult to extract while oxides of unreactive metals are most easily extracted.

• Very unreactive metals , such as gold, silver and mercury, can be obtained from their oxides by heat alone.

• Other metals from the middle of the Reactivity Series, such as zinc, iron, copper and lead, can be obtained from their oxides by heating the oxide with hydrogen, carbon (or carbon monoxide).

Iron is produced from iron ore in the blast furnace.

iron ore

iron

The Blast Furnace

• The main reactions are:• The formation of carbon monoxide

from coke (carbon):C(s) + O2 (g) CO2 (g)

C(s) + CO2 (g) 2CO(g)

• The reduction of iron oxide to iron: Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2

• Highly reactive metals, such as magnesium, aluminium, calcium, sodium and potassium, have to be obtained from their oxides by electrolysis.

Metals

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