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NATIONAL 5 CHEMISTRY UNI1 1 CHEMICAL CHANGES AND STRUCTURE.

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1 NATIONAL 5 CHEMISTRY UNI1 1 CHEMICAL CHANGES AND STRUCTURE

2 Contents Page SUB-TOPICS RATES OF REACTION ATOMIC STRUCTURE AND BONDING RELATED TO PROPERTIES OF MATERIALS FORMULAE AND REACTION QUANTITIES ACIDS AND BASES

3 RATES OF REACTION

4 REVISION OF FACTORS THAT SPEED UP A REACTION Learning intentions Describe using collision theory how rates of reaction are affected by temperature, surface area and concentration. Predict the effect any change to one of the factors above would have before experimental observation. Evaluate from experiment if the factors from above have indeed changed the speed of a chemical reaction in the manner described by collision theory.

5 Reminder from S3: What are the signs of a chemical reaction?

6 Signs of a chemical reaction are: 1. Change in colour (for example silver to reddish brown when iron rusts) 2. Change in temperature or energy, such as the production (exothermic) or loss (endothermic) of heat 3. Light, heat or sound is given off. 4. Change of form (for example, burning paper). 5. Formation of gases often appearing as bubbles. 6. Formation of precipitate (insoluble solid usually from two solutions when mixed)

7 Important Point to note In order for chemists to monitor the rate of a reaction we must observe a CHANGE in response to a chemical reaction occurring against TIME. REACTANTS  PRODUCTS e.g. magnesium + hydrochloric acid  magnesium chloride + water + hydrogen

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11 Factors Affecting the Rate of Reaction For chemical reactions to take place, the reactant particles must meet and collide with each other. Not all collisions result in a chemical reaction taking place - the collisions must take place with sufficient energy to break any bonds within the reactants. This energy, used to start the reactions is called activation energy.

12 Rates of Reaction Effect of concentrationIntroduction to collision theory Effect of temperature Contents of lesson Effect of surface area

13 Rates of reactions The speed of different chemical reactions varies hugely. Some reactions are very fast and others are very slow. fast very fast very slow fast slow Burning magnesium fireworks chemical weathering of rocks sodium and water rotting fruit Reaction Rate

14 Reactions occur when particles of reactants collide with enough energy to react (activation energy). Reactants state i.e. solid, liquid or gas will have to be taken into account when thinking about [collision theory]. Collision theory

15 Rates of Reaction Effect of concentration Introduction to collision theory Effect of temperature Effect of surface area Contents

16 Temperature and rate of reaction video Video

17 The effect of temperature on collisions How does temperature affect the rate of particle collision?

18 Temperature and particle collisions

19 Rates of Reaction Effect of concentration Introduction to collision theory Effect of temperature Effect of surface area Contents

20 What is concentration? When a solute (e.g. coffee, salt etc.) is dissolved in a solvent (e.g. water, alcohol etc.) and the resultant mixture is a solution Solute Solvent After they mix it is a solution

21 Concentration has the units (mols/l) which is referring to the amount of solute per unit volume of water we will look at this specifically much later in the topic. For example these examples all have the same volume of solvent but the amount of solute added is changed

22 Increasing concentration

23 Concentration and reaction rate Dilute hydrogen peroxide solution Concentrated hydrogen peroxide solution Add more solute (hydrogen peroxide) to the same volume Video

24 Concentration Reactions in solution involve dissolved particles. The more crowded (concentrated) the solution, the faster the reaction. Higher concentration  more particles  more collisions  faster reaction

25 Rates of Reaction Effect of concentration Introduction to collision theory Effect of temperature Effect of surface area Contents

26 Surface area and reaction rate video Increased surface area 5g Marble powder 20 ml 0.1 mol HCl 5g Marble chips medium size 20 ml 0.1 mol HCl 5g Marble lumps large size 20 ml 0.1 mol HCl Video

27 Surface area Molecules collide with the surface of the solid Extra surface for molecules to collide with. Reactions of solids can only take place at the surface of the solid. If we break a solid into smaller pieces we get a larger surface area and a faster reaction. Smaller particles  larger surface area  more collisions  faster reaction

28 If we grind up a solid to a powder we massively increase the surface area. We therefore massively increase the rate of any reaction Very fast Slow Surface area

29 Summary of the factors which can affect the rate of a reaction are;- – Particle size/surface area – Concentration of reactants – Temperature of reactants

30 Particle Size/Surface Area Decreasing the particle size of a solid in a reaction, increases the surface area available for contact between the reactants. Particle size is halved New surfaces available for contact If the surface area available for contact is increased, the rate of collisions will increase and hence the reaction rate will increase.

31 Concentration To increase the concentration of a solution you increase the amount of solute dissolved in the solvent. This means that the number of particles in solution increases. More particles means more collisions which leads to an increase in the rate of reaction.

32 Temperature As you increase the temperature of something you give it more energy making the particles move faster, so they collide more frequently. This increases the number of successful collisions and hence the reaction rate. Also, as the particles are now moving faster, the collisions between them have more kinetic energy. This means that more reactions have the required activation energy.

33 Particles and rates of reaction copy and complete

34 NATIONAL 5 CHEMISTRY CHEMICAL CHANGES AND STRUCTURE LESSON 2 MEASURING AND CALCULATING THE RATE OF REACTION

35 True or false?

36 Calculating the Rate of Reaction The rate of reaction can be calculated by measuring the : 1. Time taken for the reaction 2. One other variable (something that changes) e.g. Volume (cm 3 ) Mass (g) Concentration (moles per litre = moll -1 ) and Rate =Change (in variable) Time taken

37 Example 1 Use the graph below to calculate the rate of reaction for the first 20 seconds. Average rate = (change) / time 20 cm 3 / 20 s 1 cm 3 S -1

38 Example 2 Use the graph below to calculate the rate of reaction between 5.2 and 10.8 seconds. Average rate = (change) / time / ( ) 0.2 mol l -1 / 5.6 s mol l -1 s -1

39 Rates and Graphs These show the increasing amount of product or the decreasing amount of reactant. Amount of product Time Amount of reactant Time Steep gradient Fast reaction Shallow gradient Slow reaction Steep gradient Fast reaction Shallow gradient Slow reaction

40 Rate graphs and reactant concentrations Amount of product Time reactants product Reactant Concentration falls Rate of Reaction falls All product All reactant Mix of reactant And product Gradient of graph decreases

41 Calcium carbonate + hydrochloric acid  calcium chloride + carbon dioxide + water CaCO 3 + HCL  CaCl 2 + CO 2 + H 2 O Time (s)Decrease in mass (g)Concentration of CO Time (s)Volume of CO 2 (cm 3 ) Activity 1 plot either of these graphs on graph paper and work out the average rate for the first 120 seconds and then between 300 – 600 seconds

42 Rate and Time For some reactions, a colour change will indicate the end of the reaction. The only measurement carried out during this type of experiment is time. Therefore the rate equation is slightly changed from: Rate =Change (in variable) Time taken No other variable measured Rate= 1 Time taken Units =per second (s -1 ) To:

43 Concentration of sodium thiosulphate and the time taken for the cross to disappear Concentration of Potassium Iodide (KI) Mol/lTime (S)Rate (1/t) S Activity 2 - Plot a Time vs Rate Graph for this reaction and identify the relationship between concentration and rate of reaction Video

44 Rate = (1/t) S -1 Concentration of sodium thiosulphate (mol/l) Q) How long did it take for the cross to be obscured when the concentration of sodium thiosulphate was 0.7 mol / l? A) Rate = 1/t… so we must rearrange the equation to Time = 1/rate and obtain the rate from the graph 0.7 mol l = S -1 ; Time = 1/0.032  seconds

45 ATOMIC STRUCTURE AND BONDING RELATED TO PROPERTIES OF MATERIALS

46 Atoms and the Periodic Table. We can classify (arrange) elements in different ways:- naturally occurring/made by scientists solid/liquid/gas metal/non-metal

47 The Periodic Table of the Elements. The Periodic Table lists the chemical elements in increasing atomic number. The Periodic Table arranges elements with similar chemical properties in groups (vertical columns). All the elements in a group have similar chemical properties as they have the same number of outer electrons.

48 The Periodic Table of the elements is a useful way of classifying the elements. A vertical column of elements in the periodic table is called a group. The elements in the same group of the periodic table have similar chemical properties. The noble gases are a group of very unreactive elements.

49 Groups of elements have names: – Group 1 - Between groups 2 and 3 - Group 7 -Group 0 - the alkali metals the transition metals the halogensthe noble gases

50 Atomic Structure Electron configuration Summary activities Isotopes Introducing atoms Atomic number and mass number Contents

51 Discovery of atomic structure

52 Atoms – the building blocks All substances are made from very tiny particles called atoms. John Dalton had ideas about the existence of atoms about 200 years ago but only relatively recently have special microscopes (called electron microscopes) been invented that can ‘see’ atoms. The yellow blobs in this image are individual gold atoms, as seen through an electron microscope.

53 Elements – different types of atom Elements are the simplest substances. There are about 100 different elements. Each element is made up of just one particular type of atom, which is different to the atoms in any other element. Copper is an element made up of copper atoms only. Carbon is an element made up of carbon atoms only.

54 How small is an atom? Atoms are very small – they are about cm wide. N X3,000,000,000 If a football was enlarged by the same amount it would stretch from the UK to the USA! To make an atom the size of a football it would have to be enlarged by about 3,000,000,000 times.

55 The Amazing Atomic Zoom

56 Inside an atom Where are the electrons and nucleus found in an atom?

57 Atomic Structure Electron configuration Summary activities Isotopes Introducing atoms Atomic number and mass number Contents

58 Even smaller particles For some time, people thought that atoms were the smallest particles and could not be broken into anything smaller. proton neutron electron Scientists now know that atoms are actually made from even smaller subatomic particles. There are three types:

59 Where are subatomic particles found? Protons, neutrons and electrons are NOT evenly distributed in an atom. The electrons are spread out around the edge of the atom. They orbit the nucleus in layers called shells. The protons and neutrons exist in a dense core at the centre of the atom. This is called the nucleus.

60 The atom: check it out! Draw a labelled diagram of the atom showing the nucleus and labelling protons, neutrons and electrons. nucleus neutron proton electron

61 ParticleMassCharge proton neutron electron Properties of subatomic particles There are two properties of subatomic particles that are especially important: 1. Mass 2. Electrical charge The atoms of an element contain equal numbers of protons and electrons and so have no overall charge almost 0

62 How many protons? The atoms of any particular element always contain the same number of protons. For example: The number of protons in an atom is known as its atomic number or proton number. It is the smaller of the two numbers shown in most periodic tables. hydrogen atoms always contain 1 proton; carbon atoms always contain 6 protons; magnesium atoms always contain 12 protons,

63 What’s the atomic number? What are the atomic numbers of these elements? 11 sodium 26 iron 50 tin 9 fluorine

64 More about atomic number Each element has a definite and fixed number of protons. If the number of protons changes, then the atom becomes a different element. Changes in the number of particles in the nucleus (protons or neutrons) is very rare. It only takes place in nuclear processes such as: radioactive decay; nuclear bombs; nuclear reactors.

65 AtomProtonsNeutronsMass number hydrogen lithium aluminium Mass number Electrons have a mass of almost zero, which means that the mass of each atom results almost entirely from the number of protons and neutrons in the nucleus. The sum of the protons and neutrons in an atom’s nucleus is the mass number. It is the larger of the two numbers shown in most periodic tables

66 127 What’s the mass number? What is the mass number of these atoms? Mass number = number of protons + number of neutrons AtomProtonsNeutronsMass number helium copper cobalt iodine germanium

67 How many neutrons? How many neutrons are there in these atoms? Atom Mass number Atomic number Number of neutrons helium42 fluorine199 strontium8838 zirconium9140 uranium Number of neutrons = mass number - number of protons = mass number - atomic number

68 Building a nucleus

69 Atomic Structure Electron configuration Summary activities Isotopes Introducing atoms Atomic number and mass number Contents

70 AtomProtonsNeutronsElectrons helium copper iodine How many electrons? Atoms have no overall electrical charge and are neutral. This means atoms must have an equal number of protons and electrons. The number of electrons is therefore the same as the atomic number. Atomic number is defined as the number of protons rather than the number of electrons because atoms can lose or gain electrons but do not normally lose or gain protons

71 AtomProtonsNeutronsElectronsAtomic number Mass number boron potassium chromium mercury argon Calculating the number of electrons What are the missing numbers?

72 How are electrons arranged? Electrons are not evenly spread but exist in layers called shells. 3 rd shell 2 nd shell 1 st shell. The arrangement of electrons in these shells is often called the electron configuration.

73 How many electrons per shell? Each shell has a maximum number of electrons that it can hold. Electrons will fill the shells nearest the nucleus first. 3 rd shell holds a maximum of 8 electrons 2 nd shell holds a maximum of 8 electrons 1 st shell holds a maximum of 2 electrons

74 Calculating electron configurations

75 Properties of the nucleus and electrons

76 Thinly spread around the outside of the atom. Very small and light. Negatively charged. Found orbiting the nucleus in layers called shells. Able to be lost or gained in chemical reactions. Summary: the atom so far The nucleus is: Electrons are: Dense – it contains nearly all the mass of the atom in a tiny space. Made up of protons and neutrons. Positively charged because of the protons.

77 Atomic Structure Electron configuration Summary activities Isotopes Introducing atoms Atomic number and mass number Contents

78 What is an isotope? Elements consist of one type of atom, but sometimes these atoms can be slightly different. mass number is different atomic number is the same Atoms that differ in this way are called isotopes. Although atoms of the same element always have the same number of protons, they may have different numbers of neutrons.

79 Properties of isotopes The isotopes of an element are virtually identical in their chemical reactions. The uncharged neutrons make no difference to chemical properties but do affect physical properties such as melting point and density. Natural samples of elements are often a mixture of isotopes. This is because they have the same number of protons and the same number of electrons.

80 Isotopes of carbon Most naturally-occurring carbon exists as carbon-12, about 1% is carbon-13 and a much smaller amount is carbon protons 6 neutrons 7 electrons 6 protons 6 neutrons 8 electrons 6 protons 6 neutrons 6 electrons

81 Isotopes of hydrogen Hydrogen-1 makes up the vast majority of the naturally-occurring element but two other isotopes exist. hydrogendeuteriumtritium 1 proton 0 neutrons 1 electron 1 proton 1 neutron 1 electron 1 proton 2 neutrons 1 electron

82 Isotopes of chlorine About 75% of naturally-occurring chlorine is chlorine-35 and 25% is chlorine protons 18 neutrons 17 electrons 17 protons 20 neutrons 17 electrons

83 What are the particle numbers in each isotope? oxygen-16oxygen-18 protons neutrons electrons Isotopes of oxygen Almost all of naturally-occurring oxygen is oxygen-16 but about 0.2% is oxygen-18.

84 Isotopes and RAM Many elements are a mixture of isotopes. The RAM given in the periodic table takes account of this. For example, chlorine exists as two isotopes: chlorine-35 (75%) and chlorine-37 (25%). To calculate the RAM of a mixture of isotopes, multiply the percentage of each isotope by its atomic mass and add them together. = (0.75 x 35) + (0.25 x 37) = = 35.5 RAM of chlorine = (75% x 35) + (25% x 37)

85 Calculating RAM Bromine contains 50.5% bromine-79 and 49.5% bromine-81. = (0.505 x 79) + (0.495 x 81) = = = 80(the RAM is usually rounded to the nearest whole number) RAM of bromine = (50.5% x 79) + (49.5% x 81) What is the RAM of naturally-occurring bromine?

86 Summarizing atomic structure

87 Atomic structure word check

88 Atomic Structure Electron configuration Summary activities Isotopes Introducing atoms Atomic number and mass number Contents

89 Glossary (part 1)  atom – The smallest particle that can exist on its own.  atomic number – The number of protons in the nucleus of an element, also known as the proton number.  electron – Negative particle that orbits the nucleus of an atom.  element – Substance made up of only one type of atom.  isotopes – Different atoms of the same element. They have the same number of protons and electrons, but a different number of neutrons.

90 Glossary (part 2)  nucleus – The dense positive centre of an atom, made up of protons and neutrons.  neutron – A neutral particle, with a mass of 1. It is found in the nucleus of an atom.  mass number – The number of protons and neutrons in an atom.  proton – A positive particle, with a mass of 1. It is found in the nucleus of an atom.  relative atomic mass (RAM) – The mass of an element compared to the mass of 1 ⁄ 12 of the mass of carbon-12.

91 Anagrams

92 Atomic structure word search

93 Properties of subatomic particles

94 Multiple-choice quiz

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96 Bonding, Structure and Properties

97 Properties of Substances Everything around us is made from individual elements or compounds. These can be classified as follows: Substances Elements MetalsNon-metals Compounds IonicCovalent

98 Covalent compounds – Always contain non-metals. – e.g. CO 2, NH 3, PCl 5 Ionic compounds -Contains metals and non-metals. -e.g. NaCl, MgO, CaCO 3

99 Bonding Atoms take part in chemical reactions to achieve a full outer shell of electrons. The molecules or compounds formed are held together by chemical bonds.

100 Chemical Bonds  are invisible forces of attraction. There are 3 main types: 1.Covalent bonds (in covalent molecules) 2.Ionic bonds (in all ionic compounds) 3.Metallic bonds (in all metals)

101 Compounds and Molecules Covalent substances are made up of molecules. Molecules contain 2 or more non-metal atoms held together by covalent bonds. – E.g. chlorine molecule Cl

102 Diatomic Molecules  Contain 2 atoms only Carbon monoxide Hydrogen chloride Oxygen Nitrogen Hydrogen Group 7 elements

103 Forming Covalent Bonds Covalent Bond: Bond formed between two non- metal atoms by sharing a pair of electrons Hydrogen Molecule: Hydrogen atom + hydrogen atom = hydrogen gas Covalent bond: the two atoms share a pair of electrons Each hydrogen atoms now has 2 electrons in its outer shell. Therefore the shell is now full.

104 Starter Fill in the table below: ElementSymbolAtomic no. ProtonsElectronsNeutronsMass Number Calcium Oxygen Argon Nitrogen

105 Covalent Bonds

106 The diagram below shows how a molecule is held together by a covalent bond. e-e- e-e- A covalent bond is the force of attraction between the nuclei and the shared outer electrons.

107 The only type of covalent bond that has an equal share of electrons between the two atoms is when the two atoms are of the same element. e.g. H 2, O 2, Cl 2. These bonds are called pure covalent bonds. All other covalent bonds are classed as polar covalent bonds.

108 Polar Covalent Bonds The diagram of the covalent bond showed the outer electrons being shared equally between the bonding atoms. When covalent bonds are formed between two different types of atoms, the electrons that are shared between the atoms are not equally shared.

109 This is because some atoms attract electrons more strongly than others. This attraction is called electronegativity. In a molecule of water the bonding electrons are pulled closer to the oxygen atom. O HH e-e- e-e- e-e- e-e- ++ ++ --  = delta  - = slightly negative  + = slightly positive

110 Because the negative electrons are closer to the oxygen it becomes slightly negative. Molecule: Group of atoms held together by covalent bonds. FH e-e- e-e- ++ -- These types of bonds are known as POLAR COVALENT BONDS.

111 Bonding Diagrams N.B. Only outer electrons are involved in bonding. E.g. chlorine gas, Cl 2 Cl N.B. Outer electrons are grouped together in pairs.

112 Ammonia (nitrogen hydride), NH 3 N HH H

113 Starter 1.What are covalent bonds? 2.Why do molecules form covalent bonds? 3.How is the covalent bond held together? 4.What is a diatomic molecule?

114 Oxygen gas, O 2 This can also be drawn as O=O O O

115 Remember… The covalent bond is a result of two positive nuclei being held together by their common attraction for the shared pair of electrons.

116 Shapes of Molecules

117 The shape of a molecule is determined by the number of bonds. There are 4 basic shapes: 1.Tetrahedral Any molecule with the formula XY 4

118 2. Pyramidal Any molecule with the formula XY 3 3. Angular Any molecule with the formula XY 2 4. Planar H Cl Any molecule with the formula XY O HH

119 Starter 1.What are covalent bonds? 2.Why do molecules form covalent bonds? 3.How is the covalent bond held together? 4.What is a diatomic molecule?

120 Covalent Compounds There are two types of covalent compounds: 1.Discrete covalent molecules e.g. CO 2, HF, CH 4, N 2 2. Covalent networks e.g. C, SiO 2, CS 2

121 1. Discrete covalent molecules These are substances made up of discrete (individual) molecules. The bonds between the molecules are weaker than the covalent bonds within molecules. e.g. CO 2 O=C=O Weak intermolecular forces (Van der Waals) Strong covalent intramolecular bonds

122 Note: The intermolecular forces are weak BONDS (Van der Waals) between the molecules. All discrete molecules have low melting points because only weak intermolecular bonds are broken (Van der Waals)- not covalent bonds.

123 2. Covalent Networks These consist of a giant lattice of covalently bonded atoms. e.g. Carbon (diamond) Covalent bond Carbon atom

124 All covalent networks have very high melting points because strong covalent bonds must be broken. =Oxygen atoms =Silicon atoms

125 Starter Draw the bonding diagrams for the following – H2O – CF4 – NH3 – HF

126 Revision of Ions All atoms become more stable if they have a full outer shell of electrons (like the noble gases). During chemical reactions, atoms can lose electrons or gain electrons. When this happens the atoms become electrically charged particles called ions.

127 Atoms that lose electrons form positive ions Atoms that gain electrons form negative ions. Ion:Particle found when an atom loses or gains electron(s).

128 Ionic Compounds  Contain metal and non-metal ions.  Contain oppositely charged ion  E.g. Sodium chlorideNa + Cl - Magnesium sulphateMg 2+ SO 4 2- Remember ions are charged because the don’t have the same numbers of protons and electrons.

129 Ca 2+ P 3- Positive charge therefore more protons than electrons. Negative charge therefore less protons than electrons. 20 protons (atomic number) 18 electrons 15 protons (atomic number) 18 electrons

130 Forming Ions AtomIon NaNa + 2,8,12,8 AlAl 3+ 2,8,32,8 Metal atoms ALWAYS lose outer electrons to obtain a full shell when changing to ions.

131 AtomIon FF-FF- 2,72,8 SS 2- 2,8,62,8,8 Non-metals (except hydrogen) ALWAYS gain outer electrons to obtain a full shell when changing to ions.

132 Group Number Number of outer electrons Charge on ion Example of ion

133 Starter Draw the bonding diagrams for the following – H 2 O – CF 4 – NH 3 – HF

134 Forming Sodium Chloride The 2 atoms became ions: Na + Cl - 2,8 2,8,8

135 Structure of Ionic Compounds All ionic compounds are solid at room temperature with the ions being held in a fixed pattern called an IONIC LATTICE.

136 The ions cannot move and are held in place by the strong forces of attraction between the opposite charges. Ionic Bond: Electrostatic force of attraction between oppositely charged ions in a compound.

137 Aim: To investigate the conductivity of different substances. Method: Conductivity Experiment Beaker containing compound Wires Power supply Bulb Electrodes 6 V

138 Results: Name of substanceIs substance solid or in solution Does it have a metal in its name? Is the compound ionic, covalent or metallic? ConductorInsulator Sucrose (C 12 H 22 O 11 ) Sucrose solution (C 12 H 22 O 11 ) Polystyrene ([C 8 H 8 ] n ) Polystyrene in solution Sodium chloride solution Sodium chloride solid Copper sulphate solid Copper sulphate solution Zinc

139 Conclusion: Ionic solutions conduct electricity. Ionic solids and covalent compounds do not conduct electricity.

140 Metallic Bonding

141 Metals can lose their outer shell electrons to gain a stable electron arrangement. This is what happens in metals. Outer electrons are said to be ‘delocalised’ which allows them to move freely. This effectively creates positively charged metal ions, which are surrounded by these delocalised electrons.

142 Metal atoms form positively charged ions and so they attract the free moving electrons in the lattice. This attraction forms a metallic bond which is very strong. Metallic Bond: The electrostatic force of attraction between the positive ions in a metallic lattice and the delocalised electrons within the structure.

143 Aim: To investigate the conductivity of some metals and non-metals. Method: Conductivity Experiment Wires Power supply Material to be tested 6 V A Ammeter

144 Results: Name of substance Is substance a solid or liquid? Is substance a metal or a non-metal? ConductorInsulator Copper foil Iron Sulphur Magnesium Silicon Oxygen Aluminium Tin Lead Carbon (in form of graphite) Iodine

145 Conclusion: All metals are conductors of electricity. Non-metals are insulators of electricity, except carbon in the form of graphite.

146 Electrical Conductivity Summary Conductivity in Ionic Compounds Ionic compounds will conduct electricity when in solution or when molten (has been melted into a liquid) as the ions are free to move. An ionic compound in solution or as a melt is called an electrolyte. Solid ionic compounds do not conduct electricity because in the solid state the ionic lattice is rigid and the ions are not free to move.

147 Conductivity in Covalent Compounds Covalent solutions do not conduct electricity as they do not contain any charged particles. The exception to this is carbon in the form of graphite which can conduct electricity as it has delocalised electrons.

148 Conductivity in Metals All metals are conductors of electricity whether they are in solid or liquid form. This is because they have electrons moving freely through their structures Non-metals are insulators of electricity, except carbon in the form of graphite.

149 Starter 1.What type of bonding is present in the following compounds: a) sodium chlorideb) carbon dioxide c) potassium bromided) iron e) Silicon dioxidef) chlorine 2. Explain why sodium chloride will not conduct in solid form, but will conduct when dissolved in water.

150 Melting Points

151 Aim: To investigate the relationship between bonding and melting point. Method: Compound being tested

152 Results: CompoundFormulaBonding Melting Point ( o C) Shellac Silica Potassium iodide copper KI SiO 2 C 60 H 90 O 15 Cu Covalent discrete Covalent network Ionic Metallic o C 1723 o C 681 o C 1085 o C

153 Conclusion: In general  Ionic compounds  High melting points Metallic elements  High melting points Covalent networks  High melting points Covalent compounds (discrete)  Low melting points

154 Explanation of Melting Points To melt a substance, the bonding between units (molecules and ionic structures) has to be substantially weakened (or partly broken). To boil a substance, these bonds have to be broken.

155 Ionic Compounds  Example: sodium chloride  have very high melting and boiling points and are all solids at room temperature.  this is because ionic bonds need to be weakened/broken to melt or boil them.

156 Metallic Elements  Example: Iron  high melting and boiling points.  this is because metallic bonds need to be weakened/broken to melt or boil them.

157 Covalent Network Substances  Example: Silicon dioxide  have high melting and boiling points and are all solids at room temperature. This is because the structure is held together by strong covalent bonds.

158 Discrete Covalent Substances  Example: Carbon dioxide  have low melting and boiling points.  the Van der Waals forces that exist between covalent molecules are weak and when a discrete covalent substance is heated it is these weak forces that must be weakened or broken.  as these forces are weak they do not require much energy to break them and so melting and boiling points of discrete covalent substances are low.

159  N.B. the covalent bonds between atoms in a molecule are not broken during melting or boiling.  Almost all substances (except mercury) that exist as gases or liquids at room temperature, are discrete covalent substances.  Polar covalent substances have higher melting and boiling points than pure covalent substances due to the increased attraction between molecules.

160 Conclusion: Some covalent substances dissolve in water, but most dissolve in other solvents. Most ionic substance dissolve in water and this involves the lattice being broken up completely.

161 ‘Like dissolves like’ This means that  non-polar solutes will dissolve in non-polar solvents  polar solutes will dissolve in polar solvents  ionic solutes will dissolve in ionic/polar solvents

162 Ionic substances dissolve in water, but water isn’t ionic, it’s covalent. Remember- water has a special type of covalent bonding known as POLAR COVALENT. This is an unequal sharing of electrons within the bond, leaving one end slightly polar +ve and the other end slightly polar –ve. This is not an ion, but a polarity. Water is still covalent. Dissolving in water?

163 When an ionic solid (salt) is added to water, the water molecules arrange themselves round the lattice with the polar +ve end attracted to the –ve ion and the polar –ve end attracted to the +ve ion. If enough energy is present, the ionic lattice breaks and the compound dissolves.

164 If there is not enough energy then the solid doesn’t dissolve i.e. the solid is insoluble. Some covalent compounds also dissolve in water. These compounds must also have polar covalent bonding.

165 FORMULAE AND REACTION QUANTITIES

166 State Symbols State symbols are used in order to show what state a substance is in, suffixes are used; Elements/substances that are solids are often given the suffix (s) e.g. Ca (s) or Calcium (s) Elements/substances that are liquids are often given the suffix (l) e.g. Hg (l) or Mercury (l)

167 Elements/substances that are gases are often given the suffix (g) e.g. Ar (g) or Argon (g) Elements/substances that are dissolved in water are often given the suffix (aq) e.g. Br 2 (aq) or bromine (aq)

168 Chemical Symbolism

169 Chemical Formulae Chemical formula – gives elements involved in a compound and the number of each atom of an element in the compound. e.g. H 2 O: 2 hydrogen atoms and 1 oxygen atom per molecule.

170 Working Out Formulae – Using Valencies The chemical formula for a compound can always be worked out by considering the bonding. There is a shorter method which uses the combining powers (valency). Group Valency

171 Formulae Rules Nname Ssymbols Vvalency Sswap Ccancel out common factor Fformula

172 Simple Formulae 1. Nitrogen hydride namenitrogen hydrogen symbolsNH valency31 swap13 cancel out common factor 13 formula  NH 3

173 2. Silicon oxide name silicon oxygen symbolsSiO valency42 swap24 cancel out common factor 12 formula  SiO 2

174 3. Phosphorous iodide name phosphorous iodine symbolsPI valency31 swap13 cancel out common factor 13 formula  PI 3 Answer Questions  Textbook pg 57 Q1

175 Formulae from Prefixes If a formulae has prefixes in the name, you do not use formulae rules to work out the formula. Prefixes:Mono – 1 Di – 2 Tri– 3 Tetra– 4 Penta– 5 Hexa- 6

176 Examples CompoundFormula Sulphur dioxide Carbon tetrachloride Dinitrogen tetraoxide Nitrogen trihydride Phosphorus pentachloride SO 2 CCl 4 N2O4N2O4 NH 3 PCl 5 Answer Questions  Textbook pg 58 Q3

177 Transition Metal Valencies using Roman Numerals The block of metals in the centre of the Periodic Table are the Transition Metals. Transition metals can have more than one valency. Their valency is shown by using Roman Numerals, in the name of the compound. – E.g. Iron (III) oxide contains iron with a valency of 3.

178 Number Roman Numeral 1I 2II 3III 4IV 5V 6VI

179 1. Iron (III) chloride name iron chlorine symbols FeCl valency 3 1 swap 1 3 cancel out common factor 1 3 formula  FeCl 3

180 2. Zinc (II) oxide name zinc oxygen symbols ZnO valency 2 2 swap 2 2 cancel out common factor 1 1 formula  ZnO Answer Questions  Textbook pg 61 Q5

181 Compounds ending in ‘ate’ or ‘ite’ These compounds contain 2 elements and oxygen. – Magnesium sulphate – Lithium carbonate – Sodium hydroxide  special case Sulphate Carbonate Hydroxide Nitrate } These substances are known as COMPLEX ions which have more than one element present.

182 Their formula and VALENCY can be found on databook. Charge = Valency

183 1. Sodium sulphate (pg data book) name sodium sulphate symbolsNaSO 4 valency12 swap21 cancel out common factor 21 formula  Na 2 SO 4

184 2. Magnesium nitrate (data book) name magnesium nitrate symbols MgNO 3 valency21 swap12 cancel out common factor1 2 formula  Mg(NO 3 ) 2 N.B. Brackets are required here because there is more than one nitrate ion.

185 3. Ammonium carbonate (data book) name ammonium carbonate symbols NH 4 CO 3 valency 1 2 swap 2 1 cancel out common factor 2 1 formula  (NH 4 ) 2 CO 3 N.B. Brackets are required here because there is more than one ammonium ion.

186 Ionic Formulae

187 Revision of Ionic Compounds Ionic compounds contain oppositely charged ions. E.g. Sodium chloride Na + Cl - Magnesium sulphate Mg 2- SO 4 2- Ions are charged because they don’t have the same numbers of protons + electrons.

188 Ca 2+ P 3- Positive charge therefore more protons than electrons Negative charge therefore less protons than electrons 20 protons (atomic number) 18 electrons 15 protons (atomic number) 18 electrons

189 Forming ions AtomIon NaNa + 2,8,1 2,8 AlAl 3+ 2,8,3 2,8 Metals ALWAYS lose outer electrons to obtain a full shell when changing to ions.

190 AtomIon FF - 2,7 2,8 SS 2- 2,8,6 2,8,8 Non-metals (except H 2 ) ALWAYS gain outer electrons to obtain a full shell when forming ions.

191 Group number Charge on ion / Valency Valency = Charge

192 Ionic Formulae The ionic formulae of a compound is just the chemical formula with the charges on the ions shown. Chemical Formula Ionic Formula NaCl Na + Cl - MgO Mg 2+ O 2- CaF Ca 2+ (F - ) 2

193 1. Magnesium oxide namemagnesium oxygen symbols of ionsMg 2+ O 2- valency22 swap22 cancel out common factor 11 formula  Mg 2+ O 2- N.B. no brackets required as only 1 of each ion. Working out Ionic Formulae

194 2. Potassium flouride namepotassium fluoride symbols of ionsK + F - valency11 swap11 cancel out common factor 11 formula  K + F - N.B. no brackets required as only 1 of each ion.

195 3. Aluminium oxide namealuminium oxygen symbols of ionsAl 3+ O 2- valency32 swap23 cancel out common factor 23 formula  (Al 3+ ) 2 (O 2- ) 3 N.B. Brackets are required here as there are more than 1 of each ion.

196 3. Copper (II) chloride name copper chloride symbols of ionsCu2+Cl- valency 2 1 swap 1 2 cancel out common factor 1 2 formula  Cu 2+ (Cl - ) 2 N.B. Brackets are required here as there are more than 1 chloride ion.

197 1. Lithium carbonate (pg4 data book) name lithium carbonate symbols of ions Li + CO 3 2- valency12 swap21 cancel out common factor 21 formula  (Li+) 2 CO 3 2- N.B. Brackets are required here as there is more than 1 lithium ion. Group Ions Ionic Formulae

198 2. Ammonium sulphate (pg4 data book) name ammonium sulphate symbols of ions NH 4 + SO 4 2- valency12 swap21 cancel out common factor 21 formula  (NH 4 + ) 2 SO 4 2- N.B. Brackets are required here as there is more than 1 ammonium ion. Answer Questions  Textbook pg 63 Qs

199 Word equations Reactants  Products Always shown on left hand side- These are the starting materials for the reaction. Always shown on right hand side- These are the new substances formed in the reaction. Changes into We use word equations to describe what is happening in a chemical reaction.

200 1.Burning of carbon with oxygen: carbon + oxygen  carbon dioxide 2. Magnesium reacting with chlorine gas: magnesium + chloride  magnesium chloride 3. Iron reacting with oxygen to produce iron oxide Iron + oxygen  iron oxide

201 1. Write word equations for the reaction between a) magnesium + oxygen b) sodium and chlorine Answer a) magnesium + oxygen  magnesium oxide b)Sodium + chlorine  sodium chloride

202 Formula Equations This is when we write the chemical formula for the substances. e.g. 1. potassium hydroxide reacts with nitric acid to form potassium nitrate and water. This can be written as Potassium hyroxide + nitric acid  potassium nitrate + water KOH + HNO 3  KNO 3 + H 2 O

203 2. Calcium hydroxide reacts with hydrochloric acid to produce calcium chloride and water. This can be written as: Calcium hydroxide+ hydrochloric acid  calcium chloride + water Ca(OH) 2 + HCl  CaCl 2 + H 2 O

204 Chemical Formula You Must Know ! Name of diatomic element Formula HydrogenH2H2 NitrogenN2N2 OxygenO2O2 FluorineF2F2 ChlorineCl 2 BromineBr 2 IodineI2I2

205 Name of Compound Formula WaterH2OH2O Carbon dioxideCO 2 Hydrochloric acidHCl Sulphuric acidH 2 SO 4 Nitric acidHNO 3 Sodium hydroxideNaOH AmmoniaNH 3 Chemical Formula You Must Know !

206 Identifying atoms CompoundElement N o of atoms Element N o of atoms Salt (NaCl)Sodium1Chlorine1 Carbon dioxide (CO 2 )CarbonOxygen Water (H 2 O)HydrogenOxygen Methane (CH 4 )CarbonHydrogen Hydrogen sulphide (H 2 S) HydrogenSulphur Calcium carbonate (CaCO 3 ) CalciumCarbon Oxygen

207 Identifying atoms CompoundElement N o of atoms Element N o of atoms Salt (NaCl)Sodium1Chlorine1 Carbon dioxide (CO 2 )Carbon1Oxygen2 Water (H 2 O)HydrogenOxygen Methane (CH 4 )CarbonHydrogen Hydrogen sulphide (H 2 S) HydrogenSulphur Calcium carbonate (CaCO 3 ) CalciumCarbon Oxygen

208 Identifying atoms CompoundElement N o of atoms Element N o of atoms Salt (NaCl)Sodium1Chlorine1 Carbon dioxide (CO 2 )Carbon1Oxygen2 Water (H 2 O)Hydrogen2Oxygen1 Methane (CH 4 )CarbonHydrogen Hydrogen sulphide (H 2 S) HydrogenSulphur Calcium carbonate (CaCO 3 ) CalciumCarbon Oxygen

209 Identifying atoms CompoundElement N o of atoms Element N o of atoms Salt (NaCl)Sodium1Chlorine1 Carbon dioxide (CO 2 )Carbon1Oxygen2 Water (H 2 O)Hydrogen2Oxygen1 Methane (CH 4 )Carbon1Hydrogen4 Hydrogen sulphide (H 2 S) HydrogenSulphur Calcium carbonate (CaCO 3 ) CalciumCarbon Oxygen

210 Identifying atoms CompoundElement N o of atoms Element N o of atoms Salt (NaCl)Sodium1Chlorine1 Carbon dioxide (CO 2 )Carbon1Oxygen2 Water (H 2 O)Hydrogen2Oxygen1 Methane (CH 4 )Carbon1Hydrogen4 Hydrogen sulphide (H 2 S) Hydrogen2Sulphur1 Calcium carbonate (CaCO 3 ) CalciumCarbon Oxygen

211 Identifying atoms CompoundElement N o of atoms Element N o of atoms Salt (NaCl)Sodium1Chlorine1 Carbon dioxide (CO 2 )Carbon1Oxygen2 Water (H 2 O)Hydrogen2Oxygen1 Methane (CH 4 )Carbon1Hydrogen4 Hydrogen sulphide (H 2 S) HydrogenSulphur Calcium carbonate (CaCO 3 ) Calcium1Carbon1 Oxygen3

212 Balancing Formula Equations Three things are balanced in a chemical equation: 1. Atoms 2. Mass 3. Charge

213 When hydrogen burns in oxygen, water is formed. The formula equation for this reaction is: H 2 + O 2  H 2 O This equation is not balanced. The same amount of chemicals must be present on both sides of the reaction because atoms cannot be created or destroyed.

214 H 2 + O 2  H 2 O There are more oxygen atoms on the left hand side than on the right hand side.

215 Steps to balance an equation 1.Putting a number in front of formulae will multiply the number of atoms after it. e.g. H 2 + O 2  H 2 O LHS RHS 2 H atoms4 H atoms2 O atoms 2. We need to fix the LHS now. Do this by putting a 2 in front on H on the LHS.. H 2 + O 2  2H 2 O LHS RHS4 H atoms2 O atoms 2 2

216 More Balancing Equations 1. 2.

217 3. 4.

218 Ionic Equations 1.Reaction between magnesium + sulphuric acid Word equation Magnesium + sulphuric acid  magnesium sulphate + water Chemical equation Mg(s) +H 2 SO 4 (aq)  MgSO4(aq)+H 2 (g) Ionic equation Mg(s)+2H + Cl - (aq)  Mg 2+ (cl - ) 2 (aq)+H 2 (g)

219 The Mole

220 One mole of a substance is equal to the formula mass expressed in grams. It is also called the gram formula mass (GFM). In the previous example for ammonium phosphate, the formula mass was calculated to be 149. This means that one mole of ammonium phosphate has the mass of 149g.

221 The units for the mole is mol. When working out calculations involving formula masses and moles, you can use the triangle below to convert from mass to moles. m n FM

222 Here n= no. of moles m= mass FM= Formula mass m n FM Rearrange the triangle as follows: m = n x FM n = mFM = m FM n

223 Example 1 What is the mass of 2 moles of hydrogen chloride? STEP 1: write the chemical formula. Symbols : H Cl Valency : 1 1 Cross over valencies : 1 1 Simplest ratio : 1 1 Formula : HCl

224 STEP 2: calculate the gram formula mass. 1 x H = 1 x 1= 1 1 x Cl= 1 x 35.5= 35.5 So, Gram Formula mass = 36.5g The formula mass of the compound hydrogen chloride is This means that one mole of hydrogen chloride has a mass of 36.5g.

225 STEP 3: Rearrange triangle to calculate the mass. m = n x FM 1 mole NaCl  36.5g m = n x FM 2 x 36.5 = 73g

226 Example 2 Q. How many moles of copper (ll) carbonate are there in 494g of the substance? A. STEP 1: write the chemical formula. Symbols : Cu 2+ CO 3 2- Valency : 2 2 Cross over valencies : 1 1 Simplest ratio : 1 1 Formula : CuCO 3

227 STEP 2: Calculate the GFM CuCO 3 1 x Cu= 1 x 63.5= x C = 1 x 12= 12 3 x O= 3 x 16= 48 So, Gram Formula Mass = 123.5g

228 STEP 3: Rearrange triangle for n. Use formula n= m/FM So, n = m = 494 = 4 moles FM Therefore, 494g contains 4 moles of copper (ll) carbonate. m n FM

229 Example 3 (shortened answer) Q. Calculate the number of moles in 12.05g of magnesium sulphate. A.Chemical Formula  MgSO 4 MgSO 4 1xMg = 1 x 24.5 = xS = 1 x 32 = 32 4xO = 4 x 16 = 64 GFM = 120.5

230 n = m = = 0.1 moles FM Therefore, there are 0.1 moles in 12.05g in magnesium sulphate. m n FM

231 Mole Ratios CH 4 + O 2  CO 2 + H 2 O We can use the mole ratios from balanced equations to calculate: 1.Mass of REACTANT required for a reaction. 2.Mass of PRODUCT formed mole 2 moles 1 mole2 moles

232 Mole Ratio Calculations Q. What mass of oxygen is needed to react with 2.4g of carbon. A. C + O 2  CO 2 C : O 2 1 mole:1 mole 12g:32g 1g :32/12 2.4g:32/12 x 2.4 = 6.4g

233 Q. What mass of carbon dioxde is produced when 0.5 moles of propane is burned. A. C 3 H 8 + 5O 2  3CO 2 + 4H 2 O C 3 H 8 : CO 2 1 mole:3 moles 1 mole:132g 0.5 moles :0.5 x 132 = 66g CO 2 = 44g 3CO 2 = 132g

234 Q. What mass of hydrogen is burned to produce 100g of water. A. H 2 + ½ O 2  H 2 O 1 mole H 2 :1 mole H 2 O 1 mole H 2 O:1 mole H 2 18g:2g 1g:2/18g 100g: 2/18 x 100 = 11.11g H 2 = 2g H 2 O = 18g

235 ACIDS AND BASES

236 The pH Scale The pH scale is a continuous range of numbers from 1 to 14, which indicate the acidity or alkalinity of solutions. Acids have a pH of less than 7 (pH<7) Alkalis have a pH of more than 7 (pH>7) Pure water and neutral solutions have a pH equal to 7 (pH=7)

237 The pH of a solution can be measured using universal indicator or pH paper. The pH of a solid cannot be measured – it must first be dissolved in water or the pH paper dampened.

238 Non-Metal Oxides When a non-metal elements burns in oxygen the non-metal oxide is formed. Soluble non-metal oxides will dissolve in water to produce acidic solutions. e.g. CO 2, NO 2, SO 2, SO 3. Non-metal oxides which are insoluble, will not affect the pH.

239 Metal Oxides When metal elements burn in oxygen, the metal oxide is formed. e.g. Soluble metal oxides will dissolve in water to form the metal hydroxide. e.g. Metal oxides or hydroxides, which dissolve in water produce alkaline solutions.

240 Not all metal oxides are soluble. A general rule is that the oxides of the Group 1 and Group 2 metals produce alkaline solutions with water. Ammonia gas (NH 3 ), is soluble in water and will dissolve to produce an alkali. Insoluble metal oxides/hydroxides will not affect the pH of water.

241 Acids Acids are compounds which dissolve in water to form solutions which have a pH less than 7. The table below shows the most commonly used acids both in the lab and the home. Chemical NameFormulapHUses Hydrochloric acidHCl1-3Common lab acid Nitric acidHNO 3 1-3Common lab acid Sulphuric acidH 2 SO 4 1-3Common lab acid VinegarN/A4-6Common household acid Lemon juiceN/A4-6Common household acid Car battery acidN/A1-3Common household acid

242 Acid solutions are made by adding a substance to water which will lower its pH. The substance added could be a solid, a liquid or a gas. H + (aq) and Cl - (aq) H + (aq) and SO 4 2- (aq) H + (aq) and HCOO - (aq)

243 The substance added is made up of molecules – atoms of non-metal elements joined by covalent bonds. When the substance is added to water the covalent bonds break to form ions which become attached to water molecules. An acid solution is one which contains H + (aq) ions

244 Acid solutions are made by adding a substance to water which will lower its pH. The substance added could be a solid, a liquid or a gas. Ions involved H + (aq) and Cl - (aq) H + (aq) and SO 4 2- (aq) H + (aq) and HCOO - (aq)

245 When the substance is added to water the covalent bonds break to form ions which become attached to water molecules. An acid solution is one which contains H + (aq) ions

246 Electrolysis of Acids Acids can be electrolysed and the products collected. REMEMBER: electrolysis is the breaking up of an ionic compound in solution or the molten state using electricity.

247 All acids contain the H + ion. When electricity is passed through a solution of an acid the positively charged hydrogen ions are attracted to the negative electrode where they gain electrons to form hydrogen gas.

248 Alkalis Alkalis are compounds which dissolve in water to form solutions which have a pH greater than 7. Chemical NameFormulapHUses Sodium hydroxideNaOH11-14Common lab alkali Potassium hydroxideKOH11-14Common lab alkali Calcium hydroxideCa(OH) Common lab alkali Cleaning fluidsN/A8-14Common household alkali ToothpasteN/A8-11Common household alkali Indigestion tabletsN/A8-11Common household alkali

249 Water Water is a covalent compound. Covalent compounds do not conduct electricity. But we know that water can act as a conductor It is believed that in pure water, and all aqueous solutions, a reversible reaction takes place.

250 In water and aqueous solutions there is an equilibrium between hydrogen ions (H + ) and hydroxide ions (OH - ) and water molecules. However, there are more water molecules than there are ions and so water conducts only slightly. Pure water contains an equal concentration of hydrogen and hydroxide ions. H 2 O (l)  H + (aq) + OH - (aq)

251 To summarise: In water and any neutral solution, the concentration of H + ions and OH - ions is equal and this gives a pH = 7 (neutral). In an acid solution the concentration of hydrogen ions (H + (aq) ) is greater than the concentration of hydroxide ions (OH - (aq) ). In an alkaline solution the concentration of hydroxide ions (OH - (aq) ) is greater than the concentration of hydrogen ions (H + (aq) ).

252 In water and aqueous solutions there is an equilibrium between hydrogen ions (H + ) and hydroxide ions (OH - ) and water molecules. However, there are more water molecules than there are ions and so water conducts only slightly. Pure water contains an equal concentration of hydrogen and hydroxide ions. H 2 O (l)  H + (aq) + OH - (aq)

253 To summarise: In water and any neutral solution, the concentration of H + ions and OH - ions is equal and this gives a pH = 7 (neutral). In an acid solution the concentration of hydrogen ions (H + (aq) ) is greater than the concentration of hydroxide ions (OH - (aq) ). In an alkaline solution the concentration of hydroxide ions (OH - (aq) ) is greater than the concentration of hydrogen ions (H + (aq) ).

254 Diluting acids and alkalis Adding water to an acid increases the pH of the solution towards 7, making it less acidic by decreasing the H + (aq) ions concentration. (The acidity decreases and the pH value increases). Adding water to an alkali decreases the pH of the solution towards 7, making it less alkaline by decreasing the OH - (aq) ions concentration. (The alkalinity decreases and the pH value decreases).

255 Strength and Concentration Strength and concentration are terms that are often used incorrectly. They do not have the same meaning. Strength is used to describe the extent of dissociation into ions. Concentration is used to describe the number of moles of solute in a certain volume of solution.

256 Concentration The concentration of a solution is the number of moles of solute present in one litre of solution i.e. moles per litre (mol l -1 ). A 1 mol l -1 solution is made by dissolving 1 mole of the substance in 1 litre of solution.

257 so  n CV Volume of solution in litres Concentration of solution in mol l -1 Number of moles of solute

258 Sometimes, before calculating the concentration, there is a need to convert grams into moles. m n GFM Gram formula mass Number of moles of solute Mass in grams

259 Example calculation using Concentration Triangle Q. Calculate the concentration of a 500cm 3 solution that contains 0.1 moles. A. V= 500cm 3 = 0.5l n= 0.1 moles c= ? From triangle c= n/v = 0.1 / 0.5= 0.2 mol l -1 solution  Therefore the concentration is 0.2 mol l -1. n CV

260 Example calculation using both triangles Q. Calculate the mass of sodium hydroxide required to prepare 2 litres of mol l -1 solution. A. Look at the information we’re given first: V= 2 litres = 2l c= mol l -1 n= ? First we will calculate n n= c x v = x 2 = 0.03 moles

261 Now we know the number of mole, we can use the other triangle n= 0.03 moles FM= NaOH = = 40 m= ? m= n x FM = 0.03 x 40 = 1.2g  Therefore the mass of sodium hydroxide is 1.2g m n GFM

262 Reactions of Acids

263 Neutralisation When you add a base to an acid, the acid is slowly cancelled out or neutralised. The reaction of acids with alkalis is an example of a neutralisation. Neutralisation: reaction that moves the pH of an acid toward 7

264  Any substance which neutralises an acid by reacting with the hydrogen ions (H + ) of the acid to form water. Bases: metal oxides, metal hydroxides and metal carbonates. Neutralisation will move the pH of an acid up towards 7 and the pH of an alkali down toward 7. Bases

265 We will look at 4 different neutralisation reactions;

266 Reactions of Acids with Alkalis When an acid is neutralised by an alkali, a salt and water are formed. Remember: an alkali is a soluble metal oxide The general word equation to represent this reaction is; acid + alkali  salt + water

267 In the reaction of an acid with an alkali the hydrogen ions and hydroxide ions form water. Nitric + sodium  sodium + water acidhydroxide nitrate HNO 3 + NaOH  NaNO 3 + H 2 O

268 Reactions of Acids with Metal Oxides When a metal is burned the metal oxide is formed. If that metal oxide dissolves in water it is called an alkali. Metal oxides which do not dissolve in water can still neutralise an acid. Acids react with metal oxides to form a salt and water.

269 Acid + metal oxide  salt + water In the reaction of an acid with a metal oxide the hydrogen ions and the oxide ions form water.

270 Reaction of Acids and Metal Carbonates When an acid reacts with a metal carbonate a salt, water and carbon dioxide are formed.

271 The H + ions from the acid react with the CO 3 2- from the metal carbonate to form CO 2 and H 2 O. Test for CO 2 : Limewater turns milky if carbon dioxide if present.

272 Reactions of Acids and Metals Most metals react with acids to form a salt and hydrogen gas. The reaction that occurs is a neutralisation reaction. Copper, mercury, gold, silver and platinum do not react with acids.

273 In the reaction between a metal and an acid, the hydrogen ions of the acid from hydrogen molecules. Test for H 2 : Hydrogen burns with a pop!

274 Acid Rain Fossil fuels can burn to produce sulphur dioxide (SO 2 ). Nitrogen dioxide (NO 2 ) is produced by the sparking of air in car engine. Both SO 2 and NO 2 dissolve in water in the atmosphere to produce acid rain.

275 Acid rain damages… Buildings made from carbonate rock by eroding them. Structures made of iron and steel by making them corrode. Plant life by making the soil too acidic. Animal life by making the water in ponds and lakes too acidic.

276 Volumetric Titrations The concentration of acids/alkalis can be calculated from the results of volumetric titrations.

277 White tile Conical flask with alkali solution e.g. sodium hydroxide (NaOH) + universal indicator Acid solution e.g. dilute hydrochloric acid (HCl) Burette Pipette

278 P acid x V acid x C acid = P alkali x V alkali x C alkali P= powerV= volume in litres C= concentration Power of an acid is the number of H+ ions.  For HCl, p=1  For H 2 SO 4, p=2  For HNO 3, p=1 Power of an alkali is the number of OH- ions.  For NaOH, p=1  For Ca(OH) 2, p=2  For Al(OH) 3, p=3

279 Example 1 In a titration 10cm 3 of sodium hydroxide solution with a concentration of 0.2 mol l -1 was neutralised by 25 cm 3 of dilute hydrochloric acid. Calculate the concentration of the acid solution in moles per litre. P acid x V acid x C acid = P alkali x V alkali x C alkali 1x x C acid = 1 x 0.01x x C acid = C acid = / C acid = mol l -1

280 Example 2 What volume of 2M sodium hydroxide will be required to neutralise 40cm 3 of 2M sulphuric acid? P acid x V acid x C acid = P alkali x V alkali x C alkali 2 x 0.04 x 2 = 1 x V alkali x = V alkali x 2 V alkali = 0.16 / 2 V alkali = 0.08 litres V alkali = 80 cm 3

281 Example 3 What volume of 1M sulphuric acid will be needed to neutralise 2 litres of 0.1M potassium hydroxide? P acid x V acid x C acid = P alkali x V alkali x C alkali 2 x V acid x 1 = 1 x 2 x 0.1 V acid x 2 = 0.2 V alkali = 0.2 / 2 V alkali = 0.1 l V alkali = 100 cm 3

282 Everyday Examples of Neutralisation Taking antacid tablet to treat acid indigestion Using lime (CaO) to reduce acidity in soil and lochs.

283 Preparation of Salts

284 Salts Salts are ionic compounds which are produced in neutralisation reactions. They contain a positive ion, which comes from the acid (H + ) and a negative ion, which comes from the neutraliser. Salt: a substance in which the hydrogen ion of an acid has been replaced by a metal ion (or the ammonium ion).

285 Naming Salts First part of the name comes from the metal in the alkali. Second part of the name comes from the acid. Hydrochloric acid always forms a chloride salt. Sulphuric acid always forms a sulphate salt. Nitric acid always forms a nitrate salt.

286 AcidAlkaliSalt formed hydrochloric acid sodium hydroxide sulphuric acid calcium hydroxide nitric acid ammonium hydroxide Sodium chloride Calcium sulphate Ammonium nitrate

287 Spectator Ions The neutralisation of hydrochloric acid by sodium hydroxide can be shown as follows; The Na + (aq) and Cl - (aq) ions are not changed. We call these ions spectator ions as they do not take part in the reaction.

288 If we remove the spectator ions from the equation, the ionic equation for the reaction is; In the reaction of an acid with an alkali the hydrogen ions and hydroxide ions form water.

289 More Spectator Ions Omitting spectator ions... The spectator ions are SO 4 2- (aq) and Mg 2+

290 Making Soluble Salts Using Soluble Bases Soluble salts can be prepared by reacting acids with alkalis or bases. When an acid is neutralised by a soluble base (alkali) a salt and water are produced. An indicator is used to show when the acid has been neutralised. – E.g. Sodium chloride reacts with hydrochloric acid to produce sodium chloride (soluble salt) and water.

291 Using Metals and Insoluble Bases It is easy to make soluble salts from acids using metals. e.g. magnesium chloride (soluble salt) can be made using hydrochloric acid and magnesium metal powder. 1.Add metal powder to acid until no more reacts. 2.Filter off unreacted metal powder 3.Evaporate off the water to a small volume and then allow crystals to form naturally

292 The same filtration procedure can be used as above when producing a soluble salt using Acid + Metal Acid + Metal oxide Acid + Metal carbonate

293 Preparation of Insoluble Salts by Precipitation Insoluble salts can be prepared by precipitation reactions. Precipitation Reaction: the reaction of two solutions to form an insoluble product called a precipitate.

294 For example, silver iodide (insoluble salt) can be prepared by the following method; 1.dissolve a soluble silver salt in water 2.dissolve a soluble iodide salt in water 3.mix the two solutions together and a solid should form 4.filter the solution – the solid will be retained in the filter paper and the solution 5.wash the solid with water and allow to dry

295 The reaction taking place in this experiment is: The sodium nitrate formed is a soluble salt and so stays in solution. omitting spectator ions shows the ions which actually react


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