Standard Grade Chemistry

1
Standard Grade Chemistry

• This series of presentations is designed to help
  you revise for Standard Grade Chemistry.
• Click here for the menu.

2
Menu
To revise a topic, click on its name, when the
cursor appears
Chemical Reactions
The Periodic Table
Speed of Reaction
How Atoms Combine
Hydrocarbons
Fuels
Chemical Arithmetic
Properties of Substances
Reactions of Acids
Acids and Alkalis
Metals
Corrosion
Metals and Electricity
Carbohydrates
Plastics and Synthetic Fibres
Fertilisers
Click here to finish
3
Chemical Reactions
4
Chemical Reactions

• A chemical reaction involves the formation of new
  substances.
• How do we know that a chemical reaction has taken
  place?

• There is a change in appearance
• A precipitate is formed
• A gas is given off
• Energy is released or taken in

5
Chemical Reactions

• In an exothermic reaction energy is released.
• In an endothermic reaction energy is taken in.

6
Elements and Compounds

• All the substances in the world are made from
  about 100 elements, each of which has a name and
  a symbol. 

• A compound whose name ends in "ide" contains two
  elements only. 
• A compound whose name ends in "ite" or "ate"
  contains three elements, one of which is oxygen,
  ("ite" has less oxygen than "ate)

7
Solutions

•  A solution is formed when a material dissolves
  in a liquid.
• The material which dissolves is called a solute.
• The liquid is called a solvent.

• A material which can dissolve is soluble.
• A material which cannot dissolve is insoluble.
• A saturated solution is one where no more solute
  can dissolve.
• A table of solubility is found in the Data
  Booklet.

8
Chemical Reactions

•  Click to repeat this presentation
• Click to return to Menu
• Click to End

9
Speed of Reaction
10
Speed of Reaction

• The speed of a chemical reaction is increased
  when

• the size of the reacting particles is reduced
• the concentration of the reacting materials
  increases
• the temperature is increased

11
Catalysts

• A catalyst is a substance which

• speeds up the rate of a chemical reaction.
• is not used up in the reaction
• is not changed in the reaction
• is used in industry to reduce energy costs.

12
Enzymes

• Enzymes are biological catalysts, produced by
  living things.

• Enzymes are use to make
• Alcohol
• Medicines
• Yoghurt
• Washing powders

13
Speed of Reaction

• Click to repeat this presentation
• Click to return to Menu
• Click to End

14
The Periodic Table
15
The Periodic Table

• The elements are classified by arranging
  them in the Periodic Table.
• The atoms in the Periodic Table are listed in
  order of their Atomic Number.

16

• Elements in the Periodic Table can be classified
  in different ways.
• Solid, liquid or gas
• Metal or non-metal.
• Naturally occurring or man-made

17

• The vertical columns are called groups.
• All the elements in any group will show similar
  chemical properties.

18

• Group 1 elements are called the Alkali Metals.
• Group 7 elements are called the Halogens.
• Group 0 (or 8) gases are called the Noble Gases.
• The central block of the Periodic Table contains
  the Transition Metals.

19
Rutherfords Atom

• Elements are made of small particles called atoms
• In the centre of the atom is the nucleus,
  containing protons and neutrons.
• Electrons orbit around the nucleus, like planets
  around the Sun.

20
Atoms

• Most of the mass of the atom is found in the
  nucleus
• The nucleus contains positively charged protons.
• The nucleus also contains neutrons, which have no
  charge.
• Negatively charged electrons orbit around the
  nucleus.

21
Sub-Atomic Particles
Particle Charge Mass Location
Proton positive 1 a.m.u. nucleus
Neutron none 1 a.m.u. nucleus
Electron negative negligible In orbit around the nucleus
22
Atoms

• For each atom the Atomic Number is equal to the
  number of protons.
• The Mass Number is the number of protons
  neutrons.
• The number of neutrons is Mass Number minus
  Atomic Number.

23

• The atom is neutral because the positive charge
  of the nucleus is balanced by the negative charge
  of the electrons.
• Thus the number of electrons is the same as the
  number of protons. 

24
Electrons

• The first shell holds 2 electrons.

25
Electrons

• The first shell holds 2 electrons.
• The second shell holds 8 electrons

26
Electrons

• The first shell holds 2 electrons.
• The second shell holds 8 electron
• The third shell holds 8 electrons

27

• Since electrons are impossible to track down can
  also show them pear-shaped in electron pair
  clouds
• Each cloud can hold two electrons

28

• The number of outer electrons in an atom is the
  same as the number of its group in the Periodic
  Table.
• Atoms with the same number of outer electrons
  will have similar chemical properties.

29
Isotopes

• Not all atoms of the same element have the same
  mass. Most elements are mixtures of isotopes.
• Isotopes are atoms with the same number of
  protons but different numbers of neutrons.

30

• Relative Atomic Mass is the average mass number
  of an atom.
• It is not whole number because most elements
  consist of a mixture of isotopes.
• Different isotopes have different abundances.

31
Atoms and the Periodic Table

• Click to repeat this presentation
• Click to return to Menu
• Click to End

32
How Atoms Combine
33
How Atoms Combine

• Atoms react in such a way as to achieve a stable
  electron arrangement where they have a full outer
  electron shell. (Usually 8 electrons)

34
How Atoms Combine

• They are trying to achieve a Noble Gas structure.
  
• This means that they are trying to get the same
  electron arrangement as the nearest Noble Gas.

35
The Covalent Bond

• As two atoms come together the half-filled
  electron pair clouds overlap to form a new cloud.

36
The Covalent Bond

• As two atoms come together the half-filled
  electron pair clouds overlap to form a new cloud.
  The covalent bond

37
The Covalent Bond

• In a covalently bonded molecule the two atoms are
  held together because both nuclei are attracted
  to the shared pair of electrons.

38
Molecules

• Atoms are held together by bonds. 
• A covalent bond is formed between two atoms when
  they share a pair of electrons.
• Covalent bonds are formed between two non-metal
  atoms  

39
Molecules

• A molecule is a group of atoms, held together by
  covalent bonds.
• The molecular formula gives the number of atoms
  of each type in a covalent molecule.
• A diatomic molecule is one containing two atoms.

40

• We can write formulae by counting the atoms in a
  model or picture.

41

• We can write formulae by counting the atoms in a
  model or picture.

42
Diatomic molecules

• Hydrogen, nitrogen, oxygen, the halogens and
  carbon monoxide exist as diatomic molecules.

43

• We can draw diagrams to show the formation of
  diatomic molecules.

44

• We can draw diagrams to show the formation of
  diatomic molecules.

H H
45

• We can draw diagrams to show the formation of
  diatomic molecules.

H H ? H2
46
(No Transcript)
47
Cl Cl ? Cl2
48

• Some diatomic molecules involve more than one
  covalent bond.

49

• We can represent these molecules, using as a
  covalent bond.

• H2 H H
• Cl2 Cl Cl
• O2 O O
• N2 N N

50
Covalent molecules

• We can draw similar diagrams of discrete covalent
  molecules.

51

• Once again we can represent those more simply

ammonia NH3
methane CH4
52
Valency

• Valency is a number which helps us work out
  molecular formulae.
• It is the combining power of the atom.

• Valency is
• Group Number
• 8 minus Group Number
• Size of charge on ion
• Number after metals name e.g. copper(II)

53
Valency
Group 1 2 3 4 5 6 7 0/8
Valency 1 2 3 4 3 2 1 0
54
Chemical Formulae

• Using valency
• Write down symbols
• Write down valencies
• Swap over
• Divide (if possible)
• Formula

55
Chemical Formulae

• Using valency carbon and oxygen
• Write down symbols C O
• Write down valencies 4 2
• Swap over 2 4
• Divide (if possible) 1 2
• Formula CO2

56
Chemical Formulae

• Using valency calcium and chlorine
• Write down symbols Ca Cl
• Write down valencies 2 1
• Swap over 1 2
• Divide (if possible)
• Formula CaCl2

57
Chemical Formulae

• Using valency copper(II) nitrate
• Write down symbols Cu NO3
• Write down valencies 2 1
• Swap over 1 2
• Divide (if possible)
• Formula Cu(NO3)2

58
How Atoms Combine

•  Click to repeat this presentation
• Click to return to Menu
• Click to End

59
Fuels
60
Fuels

•  A fuel is a chemical which burns, releasing
  energy.
• An exothermic reaction is one in which heat is
  released.
• Combustion is the reaction of a substance with
  oxygen, in which energy is given out.

61

• The test for oxygen is that it relights a glowing
  splint.
• The two main gases in air are oxygen (about 20)
  and nitrogen (about 80).

62
Fuels

•  In any chemical reaction breaking bonds takes in
  energy while forming bonds releases energy.
• In an exothermic reaction the energy released by
  forming the bonds in the products is greater than
  the energy taken in to break bonds in the
  reactants.

63
Fossil Fuels

• Fossil fuels are coal, oil and natural gas which
  have been formed by the decay of natural
  materials which lived millions of years ago.
• Coal, oil and gas are finite resources i.e. the
  Earth has only limited quantities.
• A fuel crisis will occur when the amount of
  these fuels is no longer sufficient to supply our
  needs cheaply.

64

• Coal was formed from the decay of forests and
  vegetation which covered the earth 500-600
  million years ago.
• Layers built up until the heat and pressure
  changed the organic material to coal.
• Oil and natural gas were formed in a similar
  way, except that they probably came from marine
  plants and animals, compressed by layers of sand
  on the sea bed.

65

• Both coal and oil contain sulphur.
• When the fuels burn the sulphur produces a gas
  called sulphur dioxide.
• This causes pollution since it dissolves in water
  to form sulphuric acid (acid rain).
• Oil causes pollution problems if it is spilled in
  water because it does not dissolve in water and
  is poisonous to marine life.

66
Oil

• All substances have there own particular melting
  point and boiling point.
• Crude oil is a mixture of compounds which can be
  to split it into fractions.
• A fraction is a group of chemical compounds, all
  of which boil within the same temperature range.
  Oil can be separated into fractions by the
  process of fractional distillation.

67
Fractional Distillation of Oil
gases
(gaseous fuel)
petrol (gasoline)
(petrol)
naphtha
(chemicals)
paraffin (kerosine)
Heated oil from furnace
(aircraft fuel)
diesel
(fuel for lorries etc.)
residue
(wax, tar)
68
Oil Fractions
Name Carbon atoms per molecule Uses
Gases 1 to 4 Fuel
Petrol 4 to 9 Fuel for cars
Naphtha 8 to 14 Chemicals
Paraffin 10 to 16 Aircraft fuel
Diesel 15 to 20 Lorry fuel
Residue More than 20 Lubricating oil, tar, wax etc.
69
Oil Fractions

• Viscosity is a measure of the thickness of a
  liquid.
• Flammability is a measure of how easily the
  liquid catches fire.
• Volatility means how easy it is to turn the
  liquid into a gas.

70

• As the boiling point of a fraction increases
  then
• it will not evaporate as easily.
• it will be less flammable
• it will be more viscous (thicker).

71

• Moving through the fractions from gases to the
  residue
• The molecules present in the fraction are longer
  and heavier
• They will find it more difficult to become a gas
  i.e. they will be less easy to evaporate.

72

• Moving up the fractions from gases to the residue
  
• Since combustion involves the reaction of gas
  molecules with oxygen flammability will decrease.
  
• Increased molecular lengths mean that molecules
  become more "tangled up", so the liquid will
  become thicker (more viscous).

73
Tests

• The test for carbon dioxide is that it turns lime
  water cloudy.
• The test for water is that it turns anhydrous
  copper sulphate from white to blue.
• Hydrocarbons burn to produce carbon dioxide and
  water only.

74
To pump
Lime water (turns cloudy)
Anhydrous copper sulphate (turns blue)
Burning candle
75
Hydrocarbons

• When a hydrocarbon fuel burns to give carbon
  dioxide and water then
• The carbon in the carbon dioxide and the hydrogen
  in water must have come from the fuel.
• Crude oil is mainly made of compounds called
  hydrocarbons (i.e. made of carbon and hydrogen
  only).

76
Incomplete Combustion

• When fuels burn in a limited supply of air then
  incomplete combustion takes place and the
  poisonous gas, carbon monoxide (CO) is produced.
• Increasing the amount of air used to burn fuel
  improves efficiency and decreases pollution.

77
Other products of combustion

• Fossil fuels contain sulphur which produces
  sulphur dioxide when the fuel is burned.
• The oil industry tries to remove this sulphur
  from the fuels before selling them.

78
Nitrogen does not react well because of its
strong bonds.
If there is a high temperature the nitrogen and
oxygen will combine to make nitrogen oxides.
The experiment opposite shows how a high voltage
spark, like one provided by the spark plug or
lightning will do the same.
79
Nitrogen does not react well because of its
strong bonds.
If there is a high temperature the nitrogen and
oxygen will combine to make nitrogen oxides.
The experiment opposite shows how a high voltage
spark, like one provided by the spark plug or
lightning will do the same.
80
Atmospheric Pollution

• The sulphur and nitrogen oxides produced can
  dissolve in water, making acid rain.
• Unburnt hydrocarbons escaping from car exhausts
  can help cause the destruction of the ozone layer.

81
Reducing Pollution

• Air pollution caused by burning hydrocarbons can
  be reduced by
• using a special exhaust system a catalytic
  converter, in which metal catalysts (platinum or
  rhodium) will convert pollutants into harmless
  gases.
• altering the fuel to air ratio.

82
Fuels

• Click to repeat this presentation
• Click to return to Menu
• Click to End

83
Hydrocarbons
84
Homologous Series

•  An homologous series is a series of carbon
  compounds.
• The alkanes, the alkenes and the cycloalkanes are
  examples of homologous series.

85

• In a homologous series
• all members can be represented by a general
  formula.
• there is a gradation in physical properties.
• there is a similarity in chemical properties

86
Isomers

•  Isomers are compounds with the same molecular
  formula but different structural formulae
• For example C4H10

87
Alkanes

• Alkanes
• All members have a name ending in -ane.
• Alkanes have a general formula CnH2n2
• Alkanes are used as gaseous and liquid fuels, as
  well as wax and tar.

88
Alkanes

• Alkanes
• As we move down the alkanes the boiling point
  increases.
• This is because the molecular size increases,
  making it more difficult to change a molecule
  from liquid into gas.

89
Alkanes

• The alkanes, general formula CnH2n2

• methane CH4
• ethane C2H6
• propane C3H8
• butane C4H10
• pentane C5H12
• hexane C6H14
• heptane C7H16
• octane C8H18
• nonane C9H20
• decane C10H22

90
Cycloalkanes

• The cycloalkanes, general formula CnH2n

• cyclopropane C3H6
• cyclobutane C4H8
• cyclopentane C5H10
• cyclohexane C6H12

91
Alkenes

• The alkenes, general formula CnH2n

• ethene C2H4
• propene C3H6
• butene C4H8
• pentene C5H10
• hexene C6H12

92
Alkenes contain a carbon to carbon double bond
93
Isomers

• The alkenes and the cycloalkanes are isomers.
• They both have the same general formula CnH2n
• They have different structural formulae, as shown.

• CH3 CH CH2
• propene C3H6
• CH2 CH2
• CH2
• cyclopropane C3H6

94
Saturated and Unsaturated

• Saturated hydrocarbons contain only carbon to
  carbon single bonds
• Unsaturated hydrocarbons contain carbon to carbon
  double or triple bonds.

95
Unsaturated Hydrocarbons

• The test for unsaturation is that unsaturated
  hydrocarbons decolourise bromine water.
•  An addition reaction takes place when a carbon
  to carbon double bond breaks and other atoms add
  on.
• If hydrogen is added to an alkene then an alkane
  is formed.

96
Cracking Hydrocarbons

• Cracking of long-chain hydrocarbons produces
  smaller, more useful molecules.
• These molecules are unsaturated.

97
catalyst
mineral wool soaked in oil
gas
heat
98
Catalytic Cracking

• A catalyst lowers the temperature at which
  cracking takes place.
• Cracking produces some unsaturated hydrocarbons
  because there are not enough hydrogen atoms to
  produce completely saturated products.

99
Hydrocarbons

• Click to repeat this presentation
• Click to return to Menu
• Click to End

100
Plastics and Synthetic Fibres
101
Plastics and Synthetic Fibres

• Most plastics and synthetic (i.e. man-made)
  fibres come from oil.
• Plastics are selected for various uses, according
  to their properties e.g. lightness, durability,
  electrical and thermal insulation.
• Biodegradable means "able to rot away". Most
  plastics are not biodegradable and so cause
  environmental problems of disposal.

102
Burning plastics

• Certain plastics burn or smoulder to give
  poisonous fumes.
• All plastics can release carbon monoxide.
• P.V.C. can release hydrogen chloride
• Polyurethane releases hydrogen cyanide.

103
Thermoplastic or Thermosetting?

• A thermoplastic plastic is one which can be
  melted or reshaped (examples polythene,
  polystyrene, P.V.C.)
• A thermosetting plastic is one which cannot be
  melted and reshaped (examples bakelite in
  electrical fittings, formica in worktops)

104
Polymerisation

• A monomer is a small molecule which is able to
  join together with other, similar, small
  molecules.
• A polymer is the large molecule produced.
• This process is called polymerisation.
• Plastics and fibres (natural and synthetic) are
  examples of polymers. The making of plastics and
  synthetic fibres are examples of polymerisation.

105
Naming polymers

• Many polymers are made from the small unsaturated
  molecules, produced by the cracking of oil.
• The name of the polymer is derived from its
  monomer.

106
Naming polymers

• MONOMER POLYMER
•  ene poly(ene)
• ethene poly(ethene)
• propene poly(propene)
• styrene poly(styrene)
• chloroethene poly(chloroethene)
• tetrafluoroethe poly(tetrafluoroethene)

107
Addition Polymerisation

• The small unsaturated molecules add to each
  other by opening up their carbon to carbon
  double bonds.
• This process is called addition polymerisation.
• CH2CH2 CH2CH2 ?

108
Addition Polymerisation

• The small unsaturated molecules add to each
  other by opening up their carbon to carbon
  double bonds.
• This process is called addition polymerisation.
• CH2CH2 CH2CH2 ? -CH2-CH2-CH2-CH2-
• The repeat unit is (-CH2-)n

109
I
The ethene is attacked by an initiator (I) which
opens up the double bond
110
The ethene is attacked by an initiator (I) which
opens up the double bond
Another ethene adds on.
111
The ethene is attacked by an initiator (I) which
opens up the double bond
Then another
Another ethene adds on.
112
The ethene is attacked by an initiator (I) which
opens up the double bond
.
Then another
Another ethene adds on.
113
Repeat Units

•  You should be able to look at the structure of a
  polymer and work out the repeat unit and the
  monomer(s) from which it was formed.
• The repeat unit of an addition polymer is always
  only two carbon atoms long.

114
Plastics and Synthetic Fibres

• Click to repeat this presentation
• Click to return to Menu
• Click to End

115
Carbohydrates
116
Carbohydrates

•  Carbohydrates are important food for animals.
• Carbohydrates contain the elements carbon,
  hydrogen and oxygen.
• There are two hydrogen atoms for each oxygen atom
  in carbohydrates

117
Photosynthesis

• Photosynthesis is the process by which plants
  make carbohydrates and oxygen from carbon
  dioxide and water, using light energy.
• 6CO2 6H2O energy ? C6H12O6 602
  
• Chlorophyll (the green colour in plants) is used
  to absorb the light energy.

118
Respiration

•  Respiration is the process by which animals AND
  plants obtain the supply of energy that they
  need for growth, movement, warmth etc.
• They obtain this energy by breaking down the
  carbohydrate, glucose, using oxygen
• C6H12O6 602 ? 6CO2 6H20 energy
• Carbohydrates burn, releasing energy and
  producing carbon dioxide and water

119
The Atmosphere

•  The combination of respiration and
  photosynthesis lead to the balance of carbon
  dioxide/oxygen in the atmosphere.
• The clearing of forests with the loss of green
  plants, reduces the amount of photosynthesis
  taking place. This could alter the balance of
  the atmosphere, with a consequent danger to life
  on Earth.

120
Glucose

• Glucose is a carbohydrate
• Glucose is sweet
• Glucose dissolves well in water
• A beam of light can pass through glucose
  solution.
• Benedict's solution will give an orange
  precipitate with glucose.

121
Sucrose

• Sucrose is a carbohydrate
• Sucrose is sweet
• Sucrose dissolves well in water
• A beam of light can pass through sucrose
  solution.
• Benedict's solution will NOT give an an orange
  precipitate with sucrose.

122
Starch

• Starch is a carbohydrate
• Starch is not sweet
• Starch does not dissolve in water
• A beam of light cannot pass through starch
  solution.
• When iodine is added to starch a blue/ black
  colour is produced.

123
Testing Carbohydrates

• Benedict's solution (or Fehling's solution) gives
  a positive test (an orange colour) with glucose,
  fructose, maltose and other sugars but NOT
  sucrose.
• The pairs of carbohydrate glucose/fructose
  (C6H12O6) and sucrose/maltose (C12H22O11) are
  isomers because they both have the same molecular
  formula but different structural formulae.

124
Testing Carbohydrates

• Starch gives a blue/black colour when added to
  iodine.

125
Types of carbohydrates

• Monosaccharides are simple sugars with formula
  C6H12O6.
• Disaccharides are simple sugars with formula
  C12H22O11.
• Polysaccharides are complex sugars with formula
  (C6H10O5)n.

126
Condensation Polymerisation

• Glucose is a carbohydrate made in photosynthesis.
  
• Two glucose molecules join to form sucrose.
• This is a condensation reaction.
• 2C6H12O6? C12H22O11 H2O
• Glucose monomers polymerise to form starch.
• This is a condensation polymerisation.
• nC6H12O6 ? (C6H10O5)n nH2O

127
Hydrolysis

• Hydrolysis takes place when large molecules are
  broken down into smaller molecules by the
  addition of small molecules, such as water.
• The breakdown of sucrose and starch are examples
  of a hydrolysis reactions.

128
Digestion

• During digestion starch molecules are broken down
  by the body into smaller glucose molecules that
  can pass through the gut wall into the
  bloodstream.
• The breakdown of starch is brought about using
  acid or the enzyme amylase.
• Enzymes, such as amylase, are biological
  catalysts

129
Enzymes

• Enzymes, such as amylase, are biological
  catalysts
• An enzyme will work most efficiently within very
  specific conditions of temperature and pH.
• The further conditions are removed from the ideal
  the less efficiently the enzyme will perform.

130
Digestion

• Sucrose and starch molecules break down by the
  addition of water
• C12H22O11 H2O ? C6H12O6 C6H12O6
• sucrose glucose
  fructose
• (C6H10O5)n nH2O ? n C6H12O6
• starch
  glucose
• Monosaccharides have formula C6H12O6.
• Disaccharides have formula C12H22O11.

131
Alcohol

• Alcoholic drinks can be made from any fruit or
  vegetable source that is a source of sugars.
• The enzymes in yeast act as catalysts in the
  formation of alcohol.

132
Fermentation

• Fermentation is the breakdown of glucose to form
  carbon dioxide and alcohol
• C6H12O6 ? 2 CO2 2 C2H5OH
• The maximum concentration of alcohol that can be
  produced is limited because an increase in
  alcohol concentration limits the efficiency of
  the yeast.

133
Distillation of alcohol

• Since alcohol boils at 80oC and water boils at
  100oC distillation of an alcohol solution
  increases the alcohol concentration.
• Alcohol is a member of the alkanol family, called
  ethanol

134
Carbohydrates

• Click to repeat this presentation
• Click to return to Menu
• Click to End

135
Chemical Arithmetic
136
Formula Mass

• Formula mass is found by adding together the
  relative atomic masses of all the atoms present
  in the formula, e.g. calcium carbonate CaCO3
• Ca 40
• C 12
• O 16x3 48
• Formula Mass 100

137
Percentage Composition

• The percentage composition is found as follows.
• Find the formula mass of the compound.
• Find the fraction made up by the element
  required.
• Convert that fraction to a percentage.
• Percentage mass of element required x 100
• formula mass of compound

138
Percentage Composition

• Find the percentage of nitrogen in ammonium
  nitrate
• Formula NH4NO3
• Formula mass 80
• Nitrogen makes up 28 out of 80
• N (28/80) x100 35

139
Empirical Formula

• The empirical (simplest) formula is found as
  follows.
• Take the masses (or percentages) of each element
  present.
• Divide the mass of each element by its relative
  atomic mass.
• Convert these numbers into a simple, whole number
  ratio.

140
Empirical Formula

• Calculate the empirical formula for the compound
  which is 54 calcium, 43 oxygen and 3 hydrogen.
• Symbol Divide by RAM Ratio

54/40 1.35 43/16 2.7 3/1 3
54 40 3
1 2 2
Ca O H
Formula is CaO2H2 or Ca(OH)2
141
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
gfm
n
142
Chemical Equations

• Reactants are the materials with which are
  present at the start of the reaction and are
  changed by the reaction.
• Products are the materials produced by the
  chemical change.
• These are separated by an arrow (which means
  gives).
• Reactants ? Products

143
Chemical Equations

• Whenever we write a chemical equation we need to
  know
• what substances are present at the start
• what are the new substances formed in the
  chemical reaction.

144
Chemical Equations

• To know the chemical reactants and products means
  we can write a word equation
• Here we are naming the reactants and products.
  e.g.
• propane oxygen ? carbon dioxide water 

145
Chemical Equations

• We need to convert the word equation into
  symbols 
• C3H8 O2 ? CO2 H2O
• If we look closely at this equation we will
  realise that it is unbalanced there are
  different numbers of atoms on each side
• 3xC 8xH 2xO ? C 2xH 3xO

146
Chemical Equations

• We must write a balanced chemical equation where
  there are equal numbers of moles of each type of
  atom on both sides.
• We can balance the equation we have been working
  with.

147
Balancing Equations

• Propane has 3 carbons so
• C3H8 O2 ? 3CO2 H2O
• Propane has 8 hydrogens so
• C3H8 O2 ? 3CO2 4H2O
• To balance out the oxygens
• C3H8 5O2 ? 3CO2 4H2O
• This is a balanced chemical equation.

148
Using Chemical Equations

• The numbers we use to balance an equation are the
  actual numbers of moles which react.
• This gives us the mole relationship in the
  reaction.
• If we look at the example we have been given
• C3H8 5 O2 ? 3CO2 4H2O
• 1 mole 5 moles ? 3moles 4moles

149
Using Chemical Equations

• Since one mole is the formula weight in grams we
  can now work out the masses which react.
• C3H8 5 O2 ? 3CO2 4H2O
• 1 mole 5 moles ? 3mole 4moles
• 1x44g 5x32g ? 3x44g 4x72g
• 44g 160g ? 132g 72g
• Now by proportion we can work out any reacting
  quantities.

150
Using Chemical Equations

• How much oxygen is needed to burn 0.22g of
  propane?
• C3H8 5 O2 ? 3CO2 4H2O
• 1 mole 5 moles 3mole 4moles
•  To burn, 1 mole C3H8 needs 5 moles O2
•  44g C3H8 needs 160g O2
• 0.22g C3H8 needs 0.8g O2

151
Concentration of Solutions

• To connect volume, concentration and molarity of
  a solution use the triangle opposite.
• c concentration (m/l)
• n number of moles
• v volume (l)

n
v
c
152
Using Chemical Equations

• How much sodium carbonate would dissolve in 500ml
  of 0.5 m/l sulphuric acid?
• Na2CO3 H2SO4 ? Na2SO4 CO2 H2O
• 1 mole 1 mole ? 1mole 1 mole 1 mole
• 500ml of 0.5 m/l sulphuric acid contain
• 0.5 x 0.5 0.25 moles of acid.
• 0.25 moles of sodium carbonate react with 0.25
  moles of sulphuric acid.
• 1 mole sodium carbonate 106g
• 0.25 moles sodium carbonate 26.5g

153
Acid/Alkali Titrations

• 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
• N number of H/OH
• H acid
• OH alkali

154
Chemical Arithmetic

• Click to repeat this presentation
• Click to return to Menu
• Click to End

155
Properties of Substances
156
Conductivity

• An electric current is a flow of electrons.
• Conductors are materials which allow an electric
  current to pass through.
• Insulators are materials which do not allow an
  electric current to pass through.

157
Conductors

• Metals
• Graphite (a form of carbon - the only
  non-metallic conductor)
• Solutions of ionic metal compounds
• Molten ionic compounds

158
Conductivity

• A metal conducts because of its metallic bonding.
• In metallic bonding the outer electrons can jump
  from atom to atom, and thus move through the
  solid.

159
Ions and Conductivity

• Ions move through liquids.
• Positive ions are formed when atoms lose
  electrons.
• Negative ions are formed when at atoms gain
  electrons.

160
Ions

• Ions are charged particles
• Atoms gain or lose electrons to achieve the Noble
  Gas Structure.
• Positive ions are formed when metal atoms lose
  electrons.
• Negative ions are formed when at non-metal atoms
  gain electrons.

161
Ions
Gr. 1 2 3 4 5 6 7 0
e- to lose 1 2 3
e- to gain 3 2 1
ion 2 3 3- 2- -
162

• An ionic solution or a melt will conduct because
  its ions are free to move to the electrode of
  opposite sign.
• An ionic solid does not conduct because its ions
  are unable to move.

163
Liquid or Gas

• At room temperature
• A liquid or gaseous compound will be covalent.
• A liquid or gas contains small discrete
  molecules between which there are fairly small
  forces of attraction.

164
Solids

• At room temperature
• A solid compound can be ionic or covalent
• Solids are a result of very strong forces
  holding the particles together.
• Ionic solids consist of a lattice of oppositely
  charged ions.

165
Types of Solid

• In an ionic solid these forces are the ionic
  bonds i.e. the forces of attraction between the
  oppositely charged ions.

166

• A covalent network solid consists of a huge
  number of atoms held together by a network of
  covalent bonds.

167
Soluble in water?

• Most ionic substances are soluble in water, the
  lattice breaking, to free the ions  
• Most covalent substances are insoluble in water
  but can dissolve in other solvents. 

168

• An electrolyte is a substance which conducts when
  molten or in solution.
• While most ions are colourless, some are
  coloured. e.g.
• cobalt - pink/purple
• copper - blue
• dichromate - orange
• nickel - green
• permanganate - purple

169
Electrolysis

• Electrolysis occurs when d.c. (direct current) is
  passed through a melt, or an ionic solution.
  This changes the compound, releasing new
  substances at the electrodes.

170
Products of Electrolysis

• At the positive electrode chlorine, bromine,
  iodine or oxygen (from water) are released.
• At the negative electrode copper, silver or
  hydrogen (from water) are released.

171
Electrolysis

• Electrolysis of copper(II) chloride
• The positive copper ion moves to the negative
  electrode where
• Cu2 2e- ? Cu

172
Electrolysis

• Electrolysis of copper(II) chloride
• The negative chloride ion moves to the positive
  electrode where
• 2Cl- ? Cl2 2e-

173
Properties of Substances

•  Click to repeat this presentation
• Click to return to Menu
• Click to End

174
Acids and Alkalis
175
pH

•  pH is a continuous scale of acidity.
• Acids have a pH of less than 7
• Alkalis have a pH of more than 7.
• Water, and other neutral solutions have a pH of 7.

176
Oxides

• Non-metal oxides dissolve in water, giving
  acidic solutions.
• Metal oxides and hydroxides, which dissolve in
  water, give alkaline solutions.

177
Acid Rain

•  Acid Rain
• This has damaging effects on buildings and other
  structures, soil and plant and animal life.
• Sulphur dioxide gas dissolves in water in the
  atmosphere, producing sulphuric acid.
• Nitrogen oxides dissolves in water in the
  atmosphere, producing nitric acid.

178
Ions

• Acids and alkalis both contain ions. In water
  the concentration of ions is very low.
• The test for hydrogen is that it explodes with a
  "pop" when lit.

179
H and OH- ions

• Acids contain more H ions than water.
• Alkalis contain more OH- ions than water.
• Water, and other neutral solutions, contain equal
  numbers of H and OH- ions.

180
Dilution

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

181

•  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.

182
Acids and Alkalis

•  Click to repeat this presentation
• Click to return to Menu
• Click to End

183
Reactions of Acids
184
Neutralisation

•  Neutralisation is the reaction of an acid with a
  neutraliser.
• Neutralisers are metal oxides, hydroxides and
  carbonates.
• 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.

185
Neutralisation

• During a neutralisation reaction then the pH of
  any acid or alkali involved move nearer to 7.
• Neutralisation involves the reaction
• H OH- ? H2O
• A salt is the substance formed when the hydrogen
  ion of an acid is replaced by a metal (or
  ammonium) ion.

186
Salts
Acid Formula Salt Ion
hydrochloric HCl chloride Cl-
sulphuric H2SO4 sulphate SO42-
nitric HNO3 nitrate NO3-
carbonic H2CO3 carbonate CO32-
187
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

188
Acids and metals

• Acids react with some metals to release hydrogen.
  The hydrogen ions in the the acid form hydrogen
  molecules.
•  Acid rain will dissolve iron structures very
  slowly, since iron reacts with acid to produce
  hydrogen.

189
Salts

• When dilute hydrochloric acid reacts with acid
  then hydrogen and a metal chloride are formed.
• When dilute sulphuric acid reacts with acid then
  hydrogen and a metal sulphate are formed.

190
Acids, Bases and Alkalis

• A base is a substance which neutralises an acid.
• ACID BASE ? SALT WATER
• An alkali is a soluble base.
• ACID ALKALI ? SALT WATER

191
Precipitation

• An easy way to prepare salts is to react an acid
  with an insoluble metal oxide or metal
  carbonate. Excess can be removed from the
  reaction mixture by filtration.
•  Precipitation is the reaction in which two
  solutions react to form an insoluble salt.

192
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
gfm
n
193
Remember solutions?

• To connect volume, concentration and molarity of
  a solution use the triangle opposite.
• c concentration (m/l)
• n number of moles
• v volume (l)

n
v
c
194
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

195
Reactions of Acids

• Click to repeat this presentation
• Click to return to Menu
• Click to End

196
Metals
197
Metals

• Metals have a metallic lustre i.e. they are shiny
•  Metals conduct electricity when solid or liquid.
• The world's metal resources are finite and so we
  must recycle used metals.

198
Alloys

• An alloy is a mixture of metals, or a mixture of
  metal with non-metal. Examples
• Brass
• solder
• "stainless steel"

199
Properties of Metals

•   A metal's properties decide its uses
• electrical conductivity electric wiring
• thermal conductivity pots and pans
• malleability - shaped into many objects
• strength - made into certain objects

200
Reactions of Metals

• 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.

201
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.

202
Oxidation

• Oxidation is the loss of electrons by a reactant
  in a chemical reaction.
• When a metal reacts to form a compound it is
  oxidised.

203
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 reduced.

204
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.

205
The Reactivity Series

• K
• Na
• Ca
• Mg
• Al
• Zn
• Fe
• Sn
• Pb
• H
• Cu
• Hg
• Ag
• Au

• Metals are listed from most reactive to least
  reactive.

206
Reactivity

• The Reactivity Series (also called the
  Electrochemical Series) lists the metals in order
  of their ease of oxidation.
• The least active metals are those whose ions are
  most easily reduced.

207
Recovering Metals

• The less active metals do not react well and so
  occur uncombined in the earth's crust.
• Thus they were some of the first elements
  discovered.

208

• Ores are naturally occurring compounds of a
  metal.
• The more reactive metals are found combined in
  the earth's crust, as ores.
• The extraction of a metal from its ore is an
  example of reduction.

209

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

210
Recovering Metals

• 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).

211

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

212
Recovering Metals

• The more reactive a metal is, the more difficult
  it is to break down its compounds.
• Oxides of reactive metals are most difficult to
  break down.
• Oxides of unreactive metals are most easily
  broken down.

213
The Blast Furnace

• Iron is produced from iron ore in the blast
  furnace.
• There are two reactions
• 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

214
Metals

• Click to repeat this presentation
• Click to return to Menu
• Click to End

215
Metals and Electricity
216
Cells

• 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 voltage generated between different pairs of
  metals varies, and this gives rise to the
  Electrochemical Series.

217
Displacement

• Any metal, in an Electrochemical Series, will
  displace a metal below it from one of its
  compounds.
• This reaction will usually produce some visible
  signs.

218
Displacement

• If zinc reacts with copper sulphate solution the
  reactions are
• Cu2 2e ? Cu Reduction
• Zn ? Zn2 2e Oxidation
• Overall
• Cu2 Zn ? Cu Zn2 Redox
• By considering the metals with which acids will
  react it is possible to place hydrogen in the
  Electrochemical Series.

219
Chemical Energy in Cells

• Chemical changes can bring about the production
  of electrical energy.
• A cell or battery will run out (go flat) when
  the chemicals which produce electricity are used
  up.

220
Mains or Battery?

• Mains Electricity
• Cannot be transported
• Uses high voltages, which can be dangerous.
• Cheap to use
• Made from renewable energy sources

• Battery
• Is easily transported
• Uses low voltages, so is not dangerous.
• More expensive to use
• Made from finite energy sources

221
Metals and Electricity

•  Click to repeat this presentation
• Click to return to Menu
• Click to End

222
Corrosion
223
Corrosion

• Corrosion is a chemical reaction in which the
  surface of a metal changes from an element to a
  compound.
•  Different metals will corrode at different
  rates.
•  Corrosion results in metals forming compounds
  and so is an example of oxidation.

224
Rusting

• Rusting is the word used to describe the
  corrosion of iron.
• Rusting requires the presence of oxygen (from
  air) and water.
• Ferroxyl indicator, which turns blue in the
  presence of Fe2 ions can be used to show the
  extent of rusting.

225
Rusting

• The water must contain dissolved carbon dioxide
  or some other electrolyte.
• Salt acts as an electrolyte and so salt, spread
  on the roads in winter, increases the corrosion
  of car bodywork.
• Any other electrolyte would increase corrosion

226
Rusting

• When iron rusts the iron atom loses electrons to
  form iron(II) ions
• Fe ? Fe2 2e
• This is followed by a further loss of electrons
  to form the iron(III) ion
• Fe2 ?Fe3 e
•  The electrons lost by the iron are taken by the
  water and oxygen and used to form hydroxide
  ions
• 2H20 O2 4e ? 40H-

227
Rusting and electrons

• Iron does not rust when connected to the negative
  terminal of a battery because the electrons
  flowing onto the iron prevent it from losing
  electrons
• By using a cell with an iron nail, a carbon rod,
  an electrolyte, ferroxyl indicator and a
  centre-zero meter it is possible to show the
  formation of Fe2 at the iron nail and the
  movement of electrons away from the iron.

228
Rusting and electrons

• The iron atoms rust, losing electrons.

229
Rusting and electrons

• The iron atoms rust, losing electrons.
• The blue colour shows Fe2 has been formed.

230
Rusting and electrons

• The iron atoms rust, losing electrons.
• The blue colour shows Fe2 has been formed
• The centre-zero meter shows the movement of
  electrons from the iron nail to the carbon rod.

231
Electrons flow to iron

• When a cell is set up with iron and a metal (say
  Mg) higher in the Electrochemical Series then
  electrons flow to the iron.
• The reactions taking place are
• Mg ? Mg2 2e
• Fe2 2e ? Fe

Iron
Magnesium
232
Electrons flow from iron

• When a cell is set up with iron and a metal (say
  Cu) lower in the Electrochemical Series then
  electrons flow from the iron.
• The reactions taking place are
• Fe ? Fe2 2e
• Cu2 2e ? Cu

Iron
Copper
233
Electroplating

-

• Electroplating

234
Electroplating

-

• Electroplating
• The metal to be plated on is made the positive
  electrode.

Metal to be plated on
235
Electroplating

-

• Electroplating
• The metal to be plated on is made the positive
  electrode.
• The object to be coated is made the negative
  electrode.

Metal to be plated on
Metal object to be plated
236
Electroplating

-

• Electroplating
• The metal to be plated on is made the positive
  electrode.
• The object to be coated is made the negative
  electrode.
• The solution contains the ions of the metal to
  be plated on.

Metal to be plated on
Solution of plating ions
Metal object to be plated
237
Metal Plating

•  Galvanising occurs when steel (or iron) is
  coated with zinc.
• Tin-plating occurs when steel (or iron) is coated
  with tin.

238
Physical Protection

• Putting a barrier over the surface of a metal
  will provide physical protection against
  corrosion
• It will not allow air and water to come in
  contact with the metal.

• Painting
• Greasing
• Electroplating
• Galvanising
• Tin-plating
• Coating with plastic.

239
Sacraficial Protection

• Sacrificial protection
• If two metals are connected electrons will flow
  from the more active metal to the less active.
• The more active metal will corrode in preference
  to the less active metal.

• On the Finart-Grangemouth oil pipe bags of scrap
  magnesium are connected every 200 meters so
  magnesium corrodes sacraficially to protect the
  iron.

240
Tin-plating

• Tin-plating
• If it is scratched then the iron and tin are
  exposed. Since the iron is higher in the
  Electrochemical Series it will corrode in
  preference to the tin.
• The corrosion of the iron increases.

241
Galvanising

• Galvanising
• If it is scratched then the iron and zinc are
  exposed. Since the zinc is higher in the
  Electrochemical Series it will corrode in
  preference to the iron.
• The corrosion of the iron is prevented.

242
Corrosion

• Click to repeat this presentation
• Click to return to Menu
• Click to End

243
Fertilisers
244
Fertilisers

•  Increasing world population means that we need
  more efficient means of food production.
• Growing plants take nutrients from the soil.
  These nutrients include compounds of nitrogen,
  phosphorous and potassium.

245
Fertilisers

• Fertilisers are substances which are added to the
  soil to replace the essential elements needed
  for plant growth.
• Different plants require fertilisers containing
  different proportions of these nutrient elements.

246
Fertilisers

• Artificial fertilisers are soaked out of the soil
  by rain.
• They are carried into lakes and rivers where
  they increase the number of river plants.
• When these plants die then there is an increase
  in the bacteria which digest them leading to a
  decrease in oxygen in the water.
• This results in the death of fish.

247
Fixed nitrogen

• Certain plants have nitrifying bacteria present
  in nodules in their roots. These bacteria can
  convert atmospheric nitrogen (called free
  nitrogen) into nitrogen compounds (fixed
  nitrogen). These nitrogen compounds increase the
  fertility of the soil.
• Bacterial methods of increasing the nitrogen
  content of soil are cheaper than chemical methods.

248
Fixed nitrogen

• Recycling of nitrogen compounds into the soil is
  brought about by the decomposition of plant and
  animal protein by bacteria in the soil.
• Ammonium salts, potassium salts, nitrates and
  phosphates make good fertilisers because
• they contain some of the essential elements for
  plant growth (P, N and K).
• they are soluble and pass easily into the soil
  and up the plant's roots.

249
The Nitrogen Cycle

• All living things need nitrogen to make proteins.
• They cannot use free nitrogen from the
  atmosphere.
• They need to get fixed nitrogen in their food.
• The Nitrogen Cycle describes the place of
  nitrogen compounds in Nature.

250
The Nitrogen Cycle
Animals need nitrogen to make substances called
proteins
ANIMALS
251
The Nitrogen Cycle
They get this nitrogen by eating protein which
has been made by plants.
ANIMALS
PLANTS
252
The Nitrogen Cycle
plant protein eaten
ANIMALS
PLANTS
253
The Nitrogen Cycle
Plants absorb nitrates through their roots
nitrates
plant protein eaten
ANIMALS
PLANTS
254
The Nitrogen Cycle
nitrates
taken in by roots
plant protein eaten
ANIMALS
PLANTS
255
The Nitrogen Cycle
ammonia
Ammonia, NH3 comes from animal waste
nitrates
taken in by roots
plant protein eaten
ANIMALS
PLANTS
256
The Nitrogen Cycle
ammonia
sewage and manure
nitrates
taken in by roots
plant protein eaten
ANIMALS
PLANTS
257
The Nitrogen Cycle
ammonia
bacteria
sewage and manure
nitrates
Bacteria convert ammonia into nitrates
taken in by roots
plant protein eaten
ANIMALS
PLANTS
258
The Nitrogen Cycle
atmospheric nitrogen
ammonia
Thunderstorms make nitrates from N2
bacteria
sewage and manure
nitrates
taken in by roots
plant protein eaten
ANIMALS
PLANTS
259
The Nitrogen Cycle
atmospheric nitrogen
ammonia
thunder storms
bacteria
sewage and manure
nitrates
taken in by roots
plant protein eaten
ANIMALS
PLANTS
260
The Nitrogen Cycle
atmospheric nitroge