1. Unit 3
Acids, Bases and Metals

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Salt preparation
Metals
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3. Acids and Bases

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

6. 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.

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.

8. 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.

9. 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.

10. 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.

11.  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.

12. 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).

13. 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.

14. 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 c v

15. 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)

16. 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)

17. CH3CO2H(aq)H+(aq) + CH3CO2-(aq)
Weak Acids
Ethanoic acid, is a weak acid.
CH3CO2H(aq)H+(aq) + CH3CO2-(aq)

18. 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

19. 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)

20. NaOH(aq) Na+(aq) + OH-(aq) KOH(aq) K+(aq) + OH-(aq)
Strong Bases
Solutions of metal hydroxides are strong bases
NaOH(aq) Na+(aq) + OH-(aq)
KOH(aq) K+(aq) + OH-(aq)

21. NH3(g) + H2O(l)  NH4OH(aq) NH4OH(aq) NH4+(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)

22. 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

23. Acids and Bases Click here to repeat Acids and Bases
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24. Salt Preparation

25. 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.

26. 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.

27. In the reaction of an acid and an alkali the hydrogen ions and hydroxide ions form water.
H+ + OH-  H2O
HCl + NaOH  NaCl + H2O

28. 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

29. In the reaction of an acid and a metal carbonate the hydrogen ions and carbonate ions form carbon dioxide and water.
2H+ + CO32-  CO2 + H2O
2HNO3 + CaCO3  Ca(NO3)2 + CO2 + H2O

30. 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.

31. 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”.

32. 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.

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

34. 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+ + CO32-  H2O + CO2

35. 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 n
gfm

36. Remember solutions? n c v
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 c v

37. 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 = molarity
NH = number of H+ ions in acid
NOH =number of OH- ions in alkali
H = acid
OH = alkali

38. 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.

39. Salts Acid Formula Salt Ion hydrochloric HCl chloride Cl- sulphuric
H2SO4
sulphate
SO42-
nitric
HNO3
nitrate
NO3-
carbonic
H2CO3
carbonate
CO32-

40. 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.

41. 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.

42. 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.

43. Precipitation
Precipitation is the reaction in which two solutions react to form an insoluble salt.
The salt can then be filtered out and dried.

44. How to make a salt. Is salt soluble? No Yes Is base soluble? Titration
Neutralisation and
Evaporation
Precipitation

45. Ionic equations Normally when we write equations we do so like this:
HNO3 + NaOH  NaNO3 + H2O

46. We can change to write the equations using ions.
H+ + NO Na+ + OH-  Na+ + NO H2O

47. If we look closely we can see that some ions appear unchanged on both sides. We call these spectator ions.
H+ + NO Na+ + OH-  Na+ + NO H2O

48. We can rewrite the equation, without the spectator ions.
This shows the ions that participate in the reaction.
H+ + OH-  H2O

49. Salt Preparation Click here to repeat Salt Preparation
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50. Metals

51. 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.

52. 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.

53. The Electrochemical Series
We can measure the voltage produced by connecting different metals together to form a simple cell. V

54. 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.

55. The electrochemical series shows a “league table”.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.

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

57. Iron will displace copper from copper(II) sulphate solution.K Na Al Zn Fe Sn Pb Cu

58. Iron can displace copper from copper(II) sulphate solution.K Na Al Zn Fe Sn Pb Cu

59. CuSO4 Fe

60. Cu
Fe SO4

61. Iron cannot displace zinc from zinc(II) sulphate solution.K Na Al Zn Fe Sn Pb Cu

62. ZnSO4 Fe

63. Fe
ZnSO4

64. 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.

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

66. K Na Ca Mg Al Zn Fe Sn Pb Cu Hg Ag Au

67. Metals which react with acid, releasing hydrogenK Na Ca Mg Al Zn Fe Sn Pb Cu Hg Ag Au
Metals which react with acid, releasing hydrogen

68. Metals which react with acid, releasing hydrogenK Na Ca Mg Al Zn Fe Sn Pb Cu Hg Ag Au
Metals which react with acid, releasing hydrogen
Metals which do
not react with acid.

69. Metals which react with acid, releasing hydrogenK Na Ca Mg Al Zn Fe Sn Pb Cu Hg Ag Au
Metals which react with acid, releasing hydrogen H
Metals which do
not react with acid.

70. 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.

71. Put a copper rod
in a solution containing
copper ions.

72. Put a zinc rod
in a solution containing
zinc ions.

73. Connect the two metal rods through a voltmeter
There is no reading because the circuit is not complete. V

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

75. 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.

76. 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.

77. 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.

78. Oxidation and Reduction
OIL RIG
Oxidation Is Loss of electrons Reduction Is Gain of electrons
In a redox reaction oxidation and reduction go on together.

79. Redox
Ion-electron equations can be written for oxidation and reduction reactions.
These equations can be combined to produce redox equations.

80. Magnesium can be oxidised.
Mg  Mg e oxidation
Hydrogen ions can be reduced.
2H e  H2 reduction
Overall
2H+ + Mg  H2 + Mg2+ redox

81. In the displacement of copper by zinc the reaction is:
CuSO4 + Zn  ZnSO4 + Cu
This can be written as:
Cu e  Cu reduction
Zn  Zn e oxidation
Overall
Cu Zn  Cu + Zn2+ redox

82. 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.

83. 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.

84. 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.

85. 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.

86. 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.

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

88. 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).

89. iron ore
Iron is produced from iron ore in the blast furnace.
iron

90. 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

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

93. Metals Click here to repeat Metals Click here to return to the Menu
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