Presentation is loading. Please wait.

Presentation is loading. Please wait.

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

Similar presentations


Presentation on theme: "Standard Grade Chemistry This series of presentations is designed to help you revise for Standard Grade Chemistry. Click here for the menu."— Presentation transcript:

1

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

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

4 Chemical Reactions

5 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

6 Chemical Reactions In an exothermic reaction energy is released. In an endothermic reaction energy is taken in.

7 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“)

8 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.

9 Chemical Reactions Click to repeat this presentation Click to return to Menu Click to End

10 Speed of Reaction

11 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

12 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.

13 Enzymes Enzymes are biological catalysts, produced by living things. Enzymes are use to make: Alcohol Medicines Yoghurt Washing powders

14 Speed of Reaction Click to repeat this presentation Click to return to Menu Click to End

15 The Periodic Table

16 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.

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

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

19 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.

20 Rutherford’s 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.

21 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.

22 Sub-Atomic Particles ParticleChargeMassLocation Protonpositive1 a.m.u.nucleus Neutronnone1 a.m.u.nucleus ElectronnegativenegligibleIn orbit around the nucleus

23 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.

24 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.

25 Electrons The first shell holds 2 electrons.

26 Electrons The first shell holds 2 electrons. The second shell holds 8 electrons

27 Electrons The first shell holds 2 electrons. The second shell holds 8 electron The third shell holds 8 electrons

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

29 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.

30 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.

31 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.

32 Atoms and the Periodic Table Click to repeat this presentation Click to return to Menu Click to End

33 How Atoms Combine

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

35 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.

36 The Covalent Bond As two atoms come together the half-filled electron pair clouds overlap to form a new cloud.

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

38 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. e-e-e-e- + +

39 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

40 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.

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

42

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

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

45 + H + H

46 We can draw diagrams to show the formation of diatomic molecules. +  H + H  H 2

47 Cl + Cl +

48 Cl + Cl  Cl 2 + 

49 Some diatomic molecules involve more than one covalent bond. O2O2 N2N2

50 We can represent these molecules, using as a covalent bond. H 2 H H Cl 2 Cl Cl O 2 O O N 2 N N

51 Covalent molecules We can draw similar diagrams of discrete covalent molecules. ammonia NH 3 methane CH 4

52 Once again we can represent those more simply: H H N H H C H H H methane CH 4 ammonia NH 3

53 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)

54 Valency Group /8 Valency

55 Chemical Formulae Using valency 1.Write down symbols 2.Write down valencies 3.Swap over 4.Divide (if possible) 5.Formula

56 Chemical Formulae Using valency carbon and oxygen 1.Write down symbols C O 2.Write down valencies Swap over Divide (if possible) Formula CO 2

57 Chemical Formulae Using valency calcium and chlorine 1.Write down symbols Ca Cl 2.Write down valencies Swap over Divide (if possible) 5.Formula CaCl 2

58 Chemical Formulae Using valency copper(II) nitrate 1.Write down symbols Cu NO 3 2.Write down valencies Swap over Divide (if possible) 5.Formula Cu(NO 3 ) 2

59 How Atoms Combine Click to repeat this presentation Click to return to Menu Click to End

60 Fuels

61 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.

62 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%).

63 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.

64 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.

65 Coal was formed from the decay of forests and vegetation which covered the earth 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.

66 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.

67 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.

68 Fractional Distillation of Oil Heated oil from furnace gases petrol (gasoline) naphtha paraffin (kerosine) residue diesel (petrol) (gaseous fuel) (chemicals) (aircraft fuel) (fuel for lorries etc.) (wax, tar)

69 Oil Fractions NameCarbon atoms per molecule Uses Gases1 to 4Fuel Petrol 4 to 9Fuel for cars Naphtha8 to 14Chemicals Paraffin10 to 16Aircraft fuel Diesel15 to 20Lorry fuel ResidueMore than 20Lubricating oil, tar, wax etc.

70 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.

71 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).

72 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.

73 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).

74 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.

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

76 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).

77 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.

78 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.

79 Nitrogen does not react well because of its strong bonds. Air +- High Voltage spark 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 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. Air +- High Voltage spark Brown gas

81 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.

82 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.

83 Fuels Click to repeat this presentation Click to return to Menu Click to End

84 Hydrocarbons

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

86 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

87 Isomers Isomers are compounds with the same molecular formula but different structural formulae For example C 4 H 10 H C C C C H H H H C C C H H H H H C H H H

88 Alkanes All members have a name ending in -ane. Alkanes have a general formula C n H 2n+2 Alkanes are used as gaseous and liquid fuels, as well as wax and tar.

89 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.

90 Alkanes The alkanes, general formula C n H 2n+2 methaneCH 4 ethaneC 2 H 6 propaneC 3 H 8 butaneC 4 H 10 pentaneC 5 H 12 hexaneC 6 H 14 heptaneC 7 H 16 octaneC 8 H 18 nonaneC 9 H 20 decaneC 10 H 22

91 Cycloalkanes The cycloalkanes, general formula C n H 2n cyclopropane C 3 H 6 cyclobutane C 4 H 8 cyclopentane C 5 H 10 cyclohexane C 6 H 12

92 Alkenes The alkenes, general formula Cn H 2n ethene C 2 H 4 propene C 3 H 6 butene C 4 H 8 pentene C 5 H 10 hexene C 6 H 12

93 Alkenes contain a carbon to carbon double bond: H C C H H H ethene H H C C C H H H H propene H H H C C C C H H H H H butene C C

94 Isomers The alkenes and the cycloalkanes are isomers. They both have the same general formula Cn H 2n They have different structural formulae, as shown. CH 3 CH CH 2 propene C 3 H 6 CH 2 cyclopropane C 3 H 6

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

96 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.

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

98 catalyst mineral wool soaked in oil heat gas

99 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.

100 Hydrocarbons Click to repeat this presentation Click to return to Menu Click to End

101 Plastics and Synthetic Fibres

102 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.

103 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.

104 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)

105 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.

106 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.

107 Naming polymers MONOMERPOLYMER ***ene poly(***ene) ethene poly(ethene) propene poly(propene) styrene poly(styrene) chloroethene poly(chloroethene) tetrafluoroethe poly(tetrafluoroethene)

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. CH 2 =CH 2 + CH 2 =CH 2 

109 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. CH 2 =CH 2 + CH 2 =CH 2  -CH 2 -CH 2 -CH 2 -CH 2 - The repeat unit is (-CH 2 -) n

110 H C H The ethene is attacked by an initiator (I*) which opens up the double bond I*

111 The ethene is attacked by an initiator (I*) which opens up the double bond I H C C* H Another ethene adds on. H C H

112 The ethene is attacked by an initiator (I*) which opens up the double bond Another ethene adds on. I H C H C C* H Then another H C H

113 The ethene is attacked by an initiator (I*) which opens up the double bond Another ethene adds on. Then another I H C H C H C C* H ….

114 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.

115 Plastics and Synthetic Fibres Click to repeat this presentation Click to return to Menu Click to End

116 Carbohydrates

117 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

118 Photosynthesis Photosynthesis is the process by which plants make carbohydrates and oxygen from carbon dioxide and water, using light energy. 6CO 2 + 6H 2 O + energy  C 6 H 12 O Chlorophyll (the green colour in plants) is used to absorb the light energy.

119 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: C 6 H 12 O  6CO 2 + 6H energy Carbohydrates burn, releasing energy and producing carbon dioxide and water

120 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.

121 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.

122 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.

123 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.

124 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 (C 6 H 12 O 6 ) and sucrose/maltose (C 12 H 22 O 11 ) are isomers because they both have the same molecular formula but different structural formulae.

125 Testing Carbohydrates Starch gives a blue/black colour when added to iodine.

126 Types of carbohydrates Monosaccharides are simple sugars with formula C 6 H 12 O 6. Disaccharides are simple sugars with formula C 12 H 22 O 11. Polysaccharides are complex sugars with formula (C 6 H 10 O 5 ) n.

127 Condensation Polymerisation Glucose is a carbohydrate made in photosynthesis. Two glucose molecules join to form sucrose. This is a condensation reaction. 2C 6 H 12 O 6  C 12 H 22 O 11 + H 2 O Glucose monomers polymerise to form starch. This is a condensation polymerisation. nC 6 H 12 O 6  (C 6 H 10 O 5 ) n + nH 2 O

128 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.

129 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

130 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.

131 Digestion Sucrose and starch molecules break down by the addition of water: C 12 H 22 O 11 + H 2 O  C 6 H 12 O 6 + C 6 H 12 O 6 sucrose glucose fructose (C 6 H 10 O 5 ) n + nH 2 O  n C 6 H 12 O 6 starch glucose Monosaccharides have formula C 6 H 12 O 6. Disaccharides have formula C 12 H 22 O 11.

132 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.

133 Fermentation Fermentation is the breakdown of glucose to form carbon dioxide and alcohol C 6 H 12 O 6  2 CO C 2 H 5 OH The maximum concentration of alcohol that can be produced is limited because an increase in alcohol concentration limits the efficiency of the yeast.

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

135 Carbohydrates Click to repeat this presentation Click to return to Menu Click to End

136 Chemical Arithmetic

137 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 CaCO 3 Ca 40 C 12 O 16x3 48 Formula Mass 100

138 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

139 Percentage Composition Find the percentage of nitrogen in ammonium nitrate: Formula NH 4 NO 3 Formula mass = 80 Nitrogen makes up 28 out of 80 %N = (28/80) x100 = 35%

140 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.

141 Empirical Formula Calculate the empirical formula for the compound which is 54% calcium, 43% oxygen and 3% hydrogen. Symbol % Divide by RAM Ratio Ca O H /40 = /16 = 2.7 3/1 = Formula is CaO 2 H 2 or Ca(OH) 2

142 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

143 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

144 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.

145 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

146 Chemical Equations We need to convert the word equation into symbols: C 3 H 8 + O 2  CO 2 + H 2 O 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

147 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.

148 Balancing Equations Propane has 3 carbons so: C 3 H 8 + O 2  3CO 2 + H 2 O Propane has 8 hydrogens so: C 3 H 8 + O 2  3CO 2 + 4H 2 O To balance out the oxygens: C 3 H 8 + 5O 2  3CO 2 + 4H 2 O This is a balanced chemical equation.

149 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: C 3 H O 2  3CO 2 + 4H 2 O 1 mole + 5 moles  3moles + 4moles

150 Using Chemical Equations Since one mole is the formula weight in grams we can now work out the masses which react. C 3 H O 2  3CO 2 + 4H 2 O 1 mole 5 moles  3mole 4moles 1x44g 5x32g  3x44g 4x72g 44g 160g  132g 72g Now by proportion we can work out any reacting quantities.

151 Using Chemical Equations How much oxygen is needed to burn 0.22g of propane? C 3 H O 2  3CO 2 + 4H 2 O 1 mole 5 moles 3mole 4moles To burn, 1 mole C 3 H 8 needs 5 moles O 2 44g C 3 H 8 needs 160g O g C 3 H 8 needs 0.8g O 2

152 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 vc

153 Using Chemical Equations How much sodium carbonate would dissolve in 500ml of 0.5 m/l sulphuric acid? Na 2 CO 3 + H 2 SO 4  Na 2 SO 4 + CO 2 + H 2 O 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 moles of sodium carbonate react with 0.25 moles of sulphuric acid. 1 mole sodium carbonate = 106g 0.25 moles sodium carbonate = 26.5g

154 Acid/Alkali Titrations Work out unknown concentrations and volumes from the results of volumetric titrations. You use the equation V H M H N H = V OH M OH N OH V = volume M = molarity N = number of H/OH H = acid OH = alkali

155 Chemical Arithmetic Click to repeat this presentation Click to return to Menu Click to End

156 Properties of Substances

157 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.

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

159 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.

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

161 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.

162 Ions Gr e - to lose 123 e - to gain 321 ion

163 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.

164 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.

165 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.

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

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

168 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.

169 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

170 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.

171 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.

172 Electrolysis Electrolysis of copper(II) chloride The positive copper ion moves to the negative electrode where: Cu e -  Cu -+ Cu 2+

173 Electrolysis Electrolysis of copper(II) chloride The negative chloride ion moves to the positive electrode where: 2Cl -  Cl 2 + 2e - -+ Cl -

174 Properties of Substances Click to repeat this presentation Click to return to Menu Click to End

175 Acids and Alkalis

176 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.

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

178 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.

179 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.

180 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.

181 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.

182 When an acid (or alkali) is diluted then the number of H + (or OH - ) ions per cm 3 of solution decrease and so the acidity (or alkalinity) decrease.

183 Acids and Alkalis Click to repeat this presentation Click to return to Menu Click to End

184 Reactions of Acids

185 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.

186 Neutralisation During a neutralisation reaction then the pH of any acid or alkali involved move nearer to 7. Neutralisation involves the reaction: H + + OH -  H 2 O A salt is the substance formed when the hydrogen ion of an acid is replaced by a metal (or ammonium) ion.

187 Salts AcidFormulaSaltIon hydrochloricHClchlorideCl - sulphuricH 2 SO 4 sulphateSO 4 2- nitricHNO 3 nitrateNO 3 - carbonicH 2 CO 3 carbonateCO 3 2-

188 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 + + CO 3 2-  H 2 O + CO 2

189 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.

190 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.

191 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

192 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.

193 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

194 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 vc

195 Working out about neutralisations Work out unknown concentrations and volumes from the results of volumetric titrations. You use the equation V H M H N H = V OH M OH N OH V = volume M = molarity N H = number of H + ions in acid N OH =number of OH - ions in alkali H = acid OH = alkali

196 Reactions of Acids Click to repeat this presentation Click to return to Menu Click to End

197 Metals

198 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.

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

200 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

201 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.

202 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.

203 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.

204 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.

205 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.

206 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.

207 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.

208 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.

209 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.

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

211 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).

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

213 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.

214 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) + O 2 (g)  CO 2 (g) C(s) + CO 2 (g)  2CO(g) The reduction of iron oxide to iron: Fe 2 O 3 (s) + 3CO(g)  2Fe(s) + 3CO 2

215 Metals Click to repeat this presentation Click to return to Menu Click to End

216 Metals and Electricity

217 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.

218 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.

219 Displacement If zinc reacts with copper sulphate solution the reactions are: Cu e  Cu Reduction Zn  Zn e Oxidation Overall Cu 2+ + Zn  Cu + Zn 2+ Redox By considering the metals with which acids will react it is possible to place hydrogen in the Electrochemical Series.By considering the metals with which acids will react it is possible to place hydrogen in the Electrochemical Series.

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

221 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

222 Metals and Electricity Click to repeat this presentation Click to return to Menu Click to End

223 Corrosion

224 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.

225 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 Fe 2+ ions can be used to show the extent of rusting.

226 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

227 Rusting When iron rusts the iron atom loses electrons to form iron(II) ions: Fe  Fe e This is followed by a further loss of electrons to form the iron(III) ion: Fe 2+  Fe 3+ + e The electrons lost by the iron are taken by the water and oxygen and used to form hydroxide ions: 2H O 2 + 4e  40H -

228 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 Fe 2+ at the iron nail and the movement of electrons away from the iron.

229 Rusting and electrons The iron atoms rust, losing electrons.

230 Rusting and electrons The iron atoms rust, losing electrons. The blue colour shows Fe 2+ has been formed.

231 Rusting and electrons The iron atoms rust, losing electrons. The blue colour shows Fe 2+ has been formed The centre-zero meter shows the movement of electrons from the iron nail to the carbon rod.

232 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  Mg e Fe e  Fe IronMagnesium

233 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  Fe e Cu e  Cu Iron Copper

234 Electroplating +-

235 The metal to be plated on is made the positive electrode. +- Metal to be plated on

236 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

237 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 Metal object to be plated Solution of plating ions

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

239 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.

240 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.

241 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.

242 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.

243 Corrosion Click to repeat this presentation Click to return to Menu Click to End

244 Fertilisers

245 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.

246 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.

247 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.

248 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.

249 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.

250 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.

251 The Nitrogen Cycle ANIMALS Animals need nitrogen to make substances called proteins

252 The Nitrogen Cycle PLANTS ANIMALS They get this nitrogen by eating protein which has been made by plants.

253 The Nitrogen Cycle PLANTS ANIMALS plant protein eaten

254 The Nitrogen Cycle nitrates PLANTS ANIMALS plant protein eaten Plants absorb nitrates through their roots

255 The Nitrogen Cycle nitrates PLANTS ANIMALS plant protein eaten taken in by roots

256 The Nitrogen Cycle ammonia nitrates PLANTS ANIMALS plant protein eaten taken in by roots Ammonia, NH 3 comes from animal waste

257 The Nitrogen Cycle ammonia nitrates PLANTS ANIMALS plant protein eaten sewage and manure taken in by roots

258 The Nitrogen Cycle ammonia nitrates PLANTS ANIMALS plant protein eaten sewage and manure bacteria taken in by roots Bacteria convert ammonia into nitrates

259 The Nitrogen Cycle ammonia nitrates PLANTS ANIMALS atmospheric nitrogen plant protein eaten sewage and manure bacteria taken in by roots Thunderstorms make nitrates from N 2

260 The Nitrogen Cycle ammonia nitrates PLANTS ANIMALS atmospheric nitrogen plant protein eaten sewage and manure bacteria thunder storms taken in by roots

261 The Nitrogen Cycle ammonia nitrates PLANTS ANIMALS atmospheric nitrogen plant protein eaten sewage and manure bacteria thunder storms taken in by roots Death of living things also produces ammonia

262 The Nitrogen Cycle ammonia nitrates PLANTS ANIMALS atmospheric nitrogen plant protein eaten sewage and manure decaying animals after death bacteria compost thunder storms taken in by roots

263 The Nitrogen Cycle ammonia nitrates PLANTS ANIMALS atmospheric nitrogen plant protein eaten sewage and manure decaying animals after death bacteria compost thunder storms taken in by roots Some plants can change N 2 into nitrates

264 The Nitrogen Cycle ammonia nitrates PLANTS ANIMALS atmospheric nitrogen bacteria in legumes plant protein eaten sewage and manure decaying animals after death bacteria compost thunder storms taken in by roots

265 The Nitrogen Cycle Due to the way in which we live, food, containing nitrogen compounds, is sent from the country to the city. This means that the nitrogen compounds are not returned to the area from which they came. To replace that nitrogen, artificial fertilisers are needed. The Haber Process is used to make ammonia from free nitrogen.

266 The Nitrogen Cycle ammonia nitrates PLANTS ANIMALS atmospheric nitrogen bacteria on legumes plant protein eaten sewage and manure decaying animals after death bacteria compost Haber Process thunder storms taken in by roots

267 The Haber Process Ammonia is produced by the Haber Process. Nitrogen and hydrogen react over an iron catalyst to produce ammonia: N 2 + 3H 2  2NH 3 Since the formation of ammonia is a reversible reaction then not all of the nitrogen and hydrogen are converted into ammonia.

268 The Haber Process At low temperatures a large amount of ammonia is produced slowly. At high temperatures a smaller amount of ammonia is produced more quickly. This means that the Haber Process is carried out at a moderately high temperature to produce ammonia at the most economical rate.

269 The Haber Process Nitrogen (air) Hydrogen (Oil industry) Unreacted Nitrogen and hydrogen Ammonia Cooler Reaction Chamber With Fe catalyst

270 Ammonia Ammonia: has no colour. has a distinctive sharp smell. is highly soluble in water, producing an alkaline solution. can be converted into an ammonium compound by its reaction with acid: 2NH 3 + H 2 SO 4  (NH 4 ) 2 SO 4 ammonia sulphuric acid ammonium sulphate

271 Ammonia Ammonia can be prepared by the reaction of ammonium compound with alkali: NH 4 Cl + NaOH  NaCl + H 2 O + NH 3 This is also used as a test for the ammonium ion NH 4 +

272 Nitric Acid Nitrogen is a very unreactive gas. Nitric acid (HNO 3 ) is formed when nitrogen dioxide (NO 2 ), in the presence of air, dissolves in water. The presence of nitrogen oxides in the air means that they will dissolve in rain to produce a mildly acidic solution. This has the result that: –nitrogen compounds are added to the soil. –the acidity of the soil will be increased.

273 Nitric Acid Nitrogen dioxide is produced by the passage of a high voltage spark through air since a large amount of energy is required to break the bonds between the nitrogen atoms in the molecules. These conditions occur when –lightning passes through air –a spark passes in the spark plug of a car engine This does not provide an economic way of making nitrogen dioxide.

274 Nitric Acid Ammonia normally reacts with oxygen to give nitrogen and water. It will only produce nitrogen oxides if a platinum catalyst is used. This is a step in the Ostwald Process, by which nitric acid is produced. This can also be carried out in the laboratory.

275 The Ostwald Process Ammonia Oxygen (or) air Water Nitric acid Oxides of nitrogen Reaction Chamber (Pt catalyst)

276 The Ostwald Process This process is carried out at a moderately high temperature to allow it to proceed fairly quickly. Since the reaction is exothermic it is not necessary to continue heating it after the reaction has started since it will supply sufficient energy to continue at a reasonable rate.

277 Fertilisers Click to repeat this presentation Click to return to Menu Click to End

278 The End Hope you found the revision useful. Come back soon!!


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

Similar presentations


Ads by Google