CfE Chemistry Summary Notes

Transcription

1 CfE Chemistry Summary Notes S2/3 Pupil Name: You must keep this booklet safe as it will also be used in National 4/5 Chemistry.

2 Unit 1 Solutions, Compounds and Mixtures Elements and the Periodic Table Everything is made from elements in the Periodic Table. The Periodic Table contains around 109 different natural elements. Man-made elements are shown with a * next to them in the Periodic Table. Each element has its own symbol. E.g. Silver - Ag, Iron - Fe and Sodium - Na. Each element has its own atomic number. E.g. Aluminium is number 13. Each element is made from only ONE type of atom. E.g. Silver is only made from silver atoms. Compounds Compounds are made by chemically joining two or more different elements together. When the element sodium reacts together with the element chlorine we get the compound sodium chloride. Naming Compounds If a compound has the ending -IDE -ITE -ATE this means there are only 2 elements present. this means the compound has 2 elements PLUS oxygen. this means the compound also has 2 elements PLUS oxygen. Name of Compound Elements Present Lithium bromide Lithium Bromine Sodium nitrite Sodium Nitrogen Oxygen Calcium carbonate Calcium Carbon Oxygen 2

3 Word Equations Sodium metal can react with chlorine gas to produce the compound sodium chloride. This can be written in a shorter format called a word equation: REACTANTS Sodium + Chlorine PRODUCTS Sodium Chloride The means produces, makes or gives. Chemical Reactions All chemical reactions produce NEW substances. Signs of a chemical reaction include: 1. A GAS forming. This can be seen as fizzing or bubbling. This is often known as effervescence. 2. A COLOUR CHANGE. 3. A PRECIPITATE formed. This is when solutions react to form a solid. 4. During the reaction there is a change in ENERGY. This can be exothermic where heat is released or Endothermic reaction gets cold. Chemical reactions do NOT involve: Melting boiling evaporating Freezing condensing dissolving The above are known as physical changes. 3

4 Elements, Compounds and Mixtures Chemically Joined Easily separated Mixture Compound 4

5 Solute, Solvent and Solutions A solute is a substance which can be dissolved in a liquid. A solvent is a liquid in which a solute can be dissolved. A solution is what is produced when a solute is dissolved in a solvent. A saturated solution is one in which no more solute can be dissolved. A concentrate solution is one which contains a large amount of dissolved solute. A dilute solution is one which contains a small amount of dissolved solute. Solubility If a substance dissolves in water it can be called soluble in water. Eg salt If a substance does not dissolve in water it is called insoluble in water. Eg sand The higher the temperature of the solvent the faster the substance will dissolve. E.g. Sugar will dissolve faster in water at 80ºC than water at 25ºC The smaller the particle size of the substance the quicker the substance will dissolve. E.g. Sugar granules will dissolve quicker than sugar lumps. Some substances that do not dissolve in water can be dissolved in other solvents. Example: Nail varnish does not dissolve in water but does dissolve in the solvent acetone. 5

6 Separating Mixtures There are various methods used in the science lab to separate mixtures; 1. Filtration is used to separate a solid from a liquid. The solid collected in the filter paper is called the residue and the liquid collected is called the filtrate. Eg. Used to separate sand and water. 2. Distillation can be used to separate two different liquids. As different liquids can have different boiling points, one liquid can be boiled to produce a gas which is then condensed back into a liquid again. Eg. Used to separate alcohol and water. 3. Chromatography can be used to find out which liquids are contained within a mixture of liquids. Chromatography can separate the liquids. Eg. Used to separate the colours in ink. 6

7 Unit 2 Chemical Reactions Rate of Reaction How fast a new substance is made is called the speed/rate of reaction. Reactions take place when particles in the chemicals collide with each other. The rate of a reaction depends on: 1. The frequency or how often the collisions between the reactant particles happen. 2. The energy with which reactant particles collide. As the reaction progresses, the concentration of reactants decrease, and the concentration of the products increases. Concentration The units of concentration can be shown as M (eg. 1M) or mol/l (eg. 0.5mol/l). The higher the number, the greater the concentration of the solution. At a higher concentration, there are more particles in the same volume of space. This means that the particles are more likely to collide and therefore more likely to react. Low Concentration High concentration 7

8 Particle Size Any reaction involving a solid can only take place on the surface of the solid. If the solid is split into smaller pieces, the surface area increases. This means that there is an increased area for the reactant particles to collide with. The smaller the pieces, the larger the surface area. This means more collisions and a greater chance of reaction. Low Surface Area High Surface Area The following graph shows the difference in reaction rate. Powder reacts faster than lumps and therefore the powder graph has a steeper gradient than the lump graph. This explains why potatoes chopped into smaller pieces cook faster than larger lumps. Temperature If we increase the temperature of a reaction we give the particles more energy. This means they will move faster and therefore are more likely to collide with other particles. When the particles collide, they do so with more energy, and so the number of successful collisions increases. 8

9 Catalysts Another way we can speed up the rate of a chemical reaction is to add a catalyst. Unlike changing particle size, concentration or temperature, a catalyst is a substance we add to a reaction to speed the reaction up. An advantage of using a catalyst is that it is not used up during the reaction and can be used again. Enzymes are biological catalysts and can be found in animals and plants. E.g. Amylase, which is found in saliva. Hazard Symbols Corrosive Toxic Flammable Irritant 9

10 Unit 3 Elements, The Periodic Table and Bonding The Periodic Table The Periodic Table is split into groups which run vertically up and down the table. The Table is also split into periods which go horizontally across. Metals & Non-Metals Group Name Properties Examples 1 Alkali Metals Very reactive Sodium, Lithium 7 Halogens Reactive, used for killing bacteria Chlorine, Fluorine 8/0 Noble Gases Unreactive Neon, Xenon Middle section Transition Metals Can vary Copper, Silver The thick zig zag line shown in the diagram separates the metals and non-metals. Metals are found to the left of the line and non-metals to the right. This information is also found on page 3 of your databook. 10

11 Metals v. Non-Metals Metals Solid. (except MERCURY which is LIQUID) Conduct electricity. High melting and boiling point. Higher density. Non-Metals Can be solid, liquid or gas. Do NOT conduct. (except CARBON in the form of GRAPHITE) Lower melting and boiling point. Lower density. Metals have many uses; copper for wires, aluminium for cans and planes and iron for bridges and fences. The Atom The atom is made from 3 different subatomic particles. ELECTRON PROTON NEUTRON Particle Name Mass Charge Location PROTON 1 Positive Inside nucleus NEUTRON 1 No charge Inside nucleus ELECTRON 0 Negative Outside nucleus For every element: The Atomic Number The Mass Number = The number of protons. = The number of protons + neutrons. To find the number of neutrons in an element we use: Neutrons = Mass Number - Atomic Number 11

12 **In a NEUTRAL atom the number of protons is EQUAL to the number of electrons.** Nuclide Notation No. of Protons = 17 No of Electrons = 17 No of Neutrons = 35 17= 18 Electron Arrangements As the number of electrons increases they arrange themselves in a particular order in energy levels/shells. Energy level Maximum number of electrons or 18 12

13 The electron arrangement (Page 3 databook) shows how electrons are arranged in atoms. E.g. Sodium 2 electrons in its 1 st shell 8 electrons in its 2 nd shell 1 electron in its 3 rd shell The Periodic Table group number tells us the number of outer electrons the element has Eg. Group 3 elements have 3 electrons in the outer level. Ions Noble Gases have a stable electron arrangement. Noble Gases are like the rockstars of the Periodic Table, atoms want to be like them. In order to do that they LOSE or GAIN electrons to have the same electron arrangement. When this happens atoms form IONS, which are charged particles. 13

14 E.g. Na Na+ + e- Electron arrangement: 2,8,1 2,8 1 Sodium now has the electron arrangement of Neon. Rule: Metals lose electrons to form POSITIVE IONS. Non-metals gain electrons to form NEGATIVE IONS. The Covalent Bond Remember atoms want to have the electron arrangement of a noble gas so one way they can do this is by sharing electrons. A COVALENT bond is formed when 2 non-metal atoms share outer electrons. The covalent bond is very strong and is difficult to break. Covalent bonds can be single bond (Cl-Cl), double bond (O=O) or triple bond (N N). 14

15 Diatomic Molecules A diatomic molecule contains 2 atoms. There are SEVEN elements that exist as diatomic molecules on the periodic table. Hydrogen - H 2 Oxygen - O 2 Nitrogen - N 2 Fluorine - F 2 Chlorine - Cl 2 Bromine - Br 2 Iodine - I 2 Shapes of Molecules When atoms bond together the molecules produced can form different shapes. The four main shapes are Linear V-Shaped/Bent Pyramidal Tetrahedral Chemical Formula Chemical formula is a shorthand way of showing elements and compounds using symbols from the Periodic Table. Some covalent formula (non-metals) can be found by using the symbols given in the name of the compound. Eg. Silicon Tetrachloride - SiCl 4 Formula using prefixes If there is a prefix present in the name of the compound it tells you the number of atoms present. E.g. Carbon Dioxide the carbon does not have a prefix so there must be only one C, the oxygen has a DI prefix which means there are two O s. The finished formula is CO 2. 15

16 Prefix Number Mon or mono 1 Di 2 Tri 3 Tetra 4 Pent 5 Formula using the Crossover Rule If the name of the compound does not contain any prefixes we can use the Crossover Rule. The valency number of an element is used; this is the number of bonds an atom can form. Group /0 Valency Chemical formula is written using the following steps 1. Symbols 2. Valency number 3. Swap - crossover 4. Divide - if there is a common number* 5. Write the final formula remember not to show the 1 e.g C 1 H 4 is shown as CH 4 *The divide step is not always required. Symbols C H Valency 4 1 Swap 1 4 Divide* Formula CH 4 As the Transition Metals do not have a group number their valency is given using Roman Numerals 1 = (l), 2 = (ll), 3 = (lll), 4 = (lv), 5 = (V), 6 =(Vl) Eg. Nickel (II) bromide - Nickel has a valency of 2 Iron (III) oxide - Iron has a valency of 3 16

17 Examples Iron (II) Oxide Symbols Fe O Valency 2 2 Swap 2 2 Divide 1 1 Formula FeO Hydrogen Oxide Symbols H O Valency 1 2 Swap 2 1 Divide Formula H 2 O Carbon Tetrachloride* Symbols C Cl Valency Swap Divide Formula CCl 4 Silicon Hydride Symbols Si H Valency 4 1 Swap 1 4 Divide Formula SiH 4 Aluminium Oxide Symbols Al O Valency 3 2 Swap 2 3 Divide Formula Al 2 O 3 Nitrogen Trihydride* Symbols N H Valency Swap Divide Formula NH 3 Carbon Sulphide Symbols C S Valency 4 2 Swap 2 4 Divide 1 2 Formula CS 2 * do not use valency rules when the compound has a prefix 17

18 Unit 4 - Fuels Fossil Fuels Coal, oil and natural gas were created millions of years ago and are known as the Fossil Fuels. Coal was formed from dead plants that sank to the bottom of swampy water and over millions of years were buried with layers of mud. These layers of mud were then compressed by the pressure as more layers were added and formed coal. Oil and gas were formed in a similar way from tiny sea creatures that sank to the bottom of the sea millions of years ago. The pressure of more layers of sand caused the production of oil and natural gas. Fuels Fuels are substances which burn in oxygen to release energy. E.g Coal, oil, gas, wood, peat and sugar However, not all substances which give out energy are fuels. A battery is NOT a fuel as it does not burn. 18

19 Fractional Distillation Crude Oil is a mixture of different chemical compounds. By using fractional distillation we can separate the different compounds as each has a different boiling point. The fractional distillation tower is used to separate the mixture into groups with similar boiling points called fractions. The tower is hot at the bottom to collect fractions with a high boiling point, and cool at the top to collect fractions with a low boiling point. Fractional Distillation Top of the Tower Small molecules Low boiling point Low viscosity High flammability Easily evaporated Bottom of the Tower Large molecules High boiling point High viscosity Low flammability Difficult to evaporate 19

20 Composition of Air Approximately only 20% of the air around us is Oxygen. The majority is made up of Nitrogen. The ratio of Oxygen to Nitrogen in air is 1 : 4. Combustion When a fuel is burned in oxygen it is called combustion. This is an example of an exothermic reaction. There are 2 types of combustion. 1. Complete Combustion A plentiful supply of oxygen. FUEL + O 2 CO 2 (Carbon dioxide) + H 2 O (Water) 2. Incomplete Combustion A limited supply of oxygen. FUEL + O 2 C (Carbon/Soot) + CO (Carbon monoxide) + H 2 O **The test for oxygen is that it relights a glowing splint** Alkanes A homologous series is a set of compounds with similar chemical properties which can be represented by a general formula. Alkanes: are a subset of the set of hydrocarbons (contain only Hydrogen and Carbon) all end in the letters -ane are a homologous series with general formula CnH2n+2 contain single C-C bonds (saturated) 20

21 each alkane had a different name depending on how many carbon atoms are present. Prefix Number of Carbons meth- 1 eth- 2 prop- 3 but- 4 pent- 5 hex- 6 hept 7 oct- 8 E.g. The alkane with 3 carbons is called propane. Note: If you forget the prefixes then look at page 6 of the data booklet. The alkanes are listed in a table, in order, so you can work out the number of carbons from that. Structure of the Alkanes Name Formula Full Structural Formula Shortened Structural Formula Methane CH 4 CH 4 Ethane C 2 H 6 CH 3 CH 3 Propane C 3 H 8 CH 3 CH 2 CH 3 Butane C 4 H 10 CH 3 CH 2 CH 2 CH 3 21

22 Pollution The burning of fossil fuels causes POLLUTION! 1. Acid Rain Fossil fuels can often contain sulphur, which when burned produces sulphur dioxide SO 2. This acidic gas dissolves in clouds and falls as acid rain. Nitrogen oxides NO x are made by lighting or by spark plugs in car engines. This is another acidic gas which dissolves in clouds and falls as acid rain. 2. Global Warming When a fuel is burned in enough oxygen it produces CO 2 a Greenhouse gas. This gas is building up in the atmosphere and causing the world to heat up, also known as the Greenhouse Effect This could mean the melting of the polar ice caps with a rising of sea levels and more severe weather. 3. Transportation of Crude Oil Oil tankers can crash spilling oil and causing damage to the sea. Reducing Pollution To reduce pollution we have to use RENEWABLE ENERGY to make our electricity and run our cars and buses. HYDROELECTRIC --- uses falling water to make electricity. SOLAR --- uses energy from the sun WIND FARMS --- uses the wind to make electricity. WAVE POWER --- uses the waves in the sea to make electricity. BIO-FUELS --- from plants to help us run our cars and buses Renewable Fuels Ethanol, obtained from sugar cane, is a RENEWABLE fuel. 22

23 Ionic Introduction Unit 5 Acids & Alkalis Ionic bonding occurs between a metal and a non-metal. The metal gives electrons away and the non-metal accepts electrons. Metals form positive ions Non-metals form negative ions Ionic compounds from an ionic lattice. The ions are held together by electrostatic attraction. Ionic Formula The charge for each ion can be obtained from the Periodic Table as follows: Group Charge on ion ** ** ** Groups 4 and 0 do NOT form ions** Ionic formulae are worked out by writing the symbols or formulae for the positive and negative ions. Then, the positive and negative charges must be "balanced" (if they are not already the same) as below. Eg. sodium chloride Na + Cl - copper (II) sulphide Cu 2+ S 2- If the positive and negative ions don't have the same number of charges, we have to work out how many of the positive ions and how many of the negative ions would be needed to make the whole compound neutral. When we need to show more than one ion in a formula, we put brackets round the ion as below. Eg. sodium oxide (Na + ) 2 O 2- calcium bromide Ca 2+ (Br - ) 2 iron (III) chloride Fe 3+ (Cl - ) 3 aluminium oxide (Al 3+ ) 2 (O 2- ) 3 23

24 Examples Calcium oxide Symbols Ca O Valency 2 2 Swap 2 2 Divide 1 1 Chem Formula CaO Ionic Formula Ca 2+ O 2- Sodium fluoride Symbols Na F Valency 1 1 Swap 1 1 Divide Chem Formula NaF Ionic Formula Na + F - Magnesium chloride Symbols Mg Cl Valency 2 1 Swap 1 2 Divide Chem Formula MgCl 2 Ionic Formula Mg 2+ (Cl - ) 2 Potassium sulphide Symbols K S Valency 1 2 Swap 2 1 Divide Chem Formula K 2 S Ionic Formula (K + ) 2 S 2- Formula involving complex ions Complex ions contain more than one kind of atom. These are found on pg 6 of the databook. E.g. CO 3 2-, NO 3 -, NH 4 + and OH - The valency of these ions is the same as their charge. The formula is worked out in the same way, using the crossover method. Examples Potassium nitrate Ammonium chloride Symbols K NO 3 Valency 1 1 Swap 1 1 Divide Chem Formula KNO 3 Ionic Formula K + - NO 3 Symbols NH 4 Cl Valency 1 1 Swap 1 1 Divide Chem Formula NH 4 Cl Ionic Formula NH + 4 Cl - 24

25 Neutral Calcium hydroxide Sodium carbonate Symbols Ca OH Valency 2 1 Swap 1 2 Divide Chem Formula Ca(OH) 2 Ionic Formula Ca 2+ (OH - ) 2 Symbols Na CO 3 Valency 1 2 Swap 2 1 Divide Chem Formula Na 2 CO 3 Ionic Formula (Na + 2- ) 2 CO 3 The ph Scale ph Scale measures how acidic or alkaline a solution is. The ph can be found using universal indicator, litmus paper or using a ph meter. Acid Alkali Common lab Hydrochloric acid HCl Sulphuric acid H 2 SO 4 Sodium hydroxide NaOH Nitric acid HNO 3 Potassium hydroxide KOH Common household Coke lemonade Bleach Oven cleaner All acids contain the hydrogen ion H + All alkalis contain the hydroxide ion OH High conc. H + ions Acid s Low conc. H + ions Low conc. OH - ions Alkalis High conc. OH - ions Acids ph less than 7 Neutral ph 7 Alkali - ph more than 7 25

26 Making Acids When we burn certain non-metals in the presence of O 2, a non-metal oxide is produced. State Symbols Solid s Liquid l Gas g Aqueous aq E.g. C(s) + O 2 (g) CO 2 (g) The non-metal oxide, if soluble, can then be dissolved in water to produce an acid. CO 2 (g) + H 2 O(l) (H + ) 2 CO 3 2- (aq) (Carbonic acid) Acid Rain Coal contains sulphur which when burned produces sulphur dioxide. S(s) + O 2 (g) SO 2 (g) When the sulphur dioxide rises and is absorbed by the clouds sulphuric acid is formed. SO 2 (g) + H 2 O(l) (H + ) 2 SO 4 2- (aq) (Sulphuric acid) Other causes of acid rain The energy from lightening in thunderstorms and the energy from the spark plug in an engine allows the nitrogen and oxygen in air to react together. The nitrogen dioxide is then absorbed into the clouds to form nitric acid. NO 2 (g) + H 2 O(l) H + NO 3 - (aq) (Nitric Acid) Making Alkalis If a metal oxide or metal hydroxide dissolves in water, an alkali is formed. i) Sodium oxide dissolves to form sodium hydroxide solution. ii) Calcium hydroxide dissolves to form calcium hydroxide solution. 26

27 Diluting Acids & Alkalis When water is added to an acid the ph moves towards 7 as the concentration of H + ions decreases. When water is added to an alkali the ph moves towards 7 as the concentration of OH - decreases. Neutralisation Acid & Alkali Acids can be neutralised by certain substances to make neutral compounds (ph 7). The reaction is called neutralisation and the substance which reacts with the acid is called a neutraliser or a base. During neutralisation a SALT is always produced. Everyday neutralisation reactions include: a) adding lime to rivers and lochs to reduce the effects of acid rain. b) treating a bee sting (acid) with baking soda (alkaline). c) treating a wasp sting (alkaline) with vinegar (acid). 27

28 Neutralisation Reactions When a salt is produced in a neutralisation reaction part of the name comes from the acid used in the reaction and the other part from the neutraliser/base. Acid Hydrochloric Sulphuric Nitric Salt Produced Chloride Sulphate Nitrate 1. Acid & Alkali Acid + Alkali (neutraliser) Salt (neutral) + Water (neutral) The hydrogen (H + ) from the acid reacts with the hydroxide (OH - ) from the alkali to form water. Eg. Full equation H + Cl - (aq) + NaOH - (aq) Na + Cl - + H 2 O(l) H + (aq) + OH - (aq) H 2 O(l) (ph 7) 2. Acid & Reactive metal Reactive Metal + Acid Salt + Hydrogen Eg. Magnesium + nitric acid Magnesium nitrate + Hydrogen **The test for hydrogen gas is that is burns with a pop.** 3. Acid & Carbonate Carbonate + Acid Salt + Carbon dioxide + Water Eg. Calcium carbonate + Hydrochloric acid Calcium chloride + Carbon dioxide + Water **The test for carbon dioxide is that it turns limewater cloudy/milky** 28

29 The Mole The mole is the formula mass of an element, compound or molecule, expressed in grams. 1. For Elements The gram formula mass of Aluminium is 27 grams, this information is found in page 6 of your databook. Therefore 1 Mole of Aluminium is 27 grams. Note: For diatomic elements the databook value is doubled. E.g. for chlorine Cl 2 the mass of one mole is 71g not 35.5g. 2. For Molecules or Compounds The chemical formula of a molecule or compound tells us which elements are present, and in what quantity. Calculating the mass of all the elements present will give the mass of one mole of the molecule or compound. E.g. Lithium Oxide (Li 2 O). The formula tells us there are 2 atoms of lithium and 1 atom of oxygen present. Therefore: 1 Mole of Li 2 O = 30g E.g. Sodium Hydroxide (NaOH) The formula tells us there is 1 atom of sodium, 1 atom of oxygen and 1 atom of hydrogen present. 1 Mole of NaOH = 40g 29

30 Unit 6 Metals, Reactivity Series, Electricity & Corrosion Properties of metals Iron is a metal that is strong and so it can be used to build bridges and railway tracks. Copper metal has many uses. It has good electrical conduction. This means that it is very good at conducting electricity so it is used to make electrical wires. Copper has good thermal conductivity. This means that it is good at letting heat move through it and so it is sometimes used to make cooking pots. Copper does not corrode so it is also good for making water pipes. Aluminium is a metal that has a very low density. This means that it is light so it is used to make aeroplanes. Metals that are used to make jewellery, have to be malleable. This means that they are can be hammered easily into different shapes. Gold, silver and platinum are very malleable metals. These metals are also used because they do not corrode and stay shiny for a long time. Tin metal does not corrode so it is a good metal to use to make food cans. Alloys Pure metals do not always have the properties that we need for a particular job. An alloy is a substance made by melting and mixing metals with other elements. Alloys are often more useful than pure metals because they have different properties from the pure metals. This can make them more suitable for certain uses. E.g. Solder is an alloy of TIN and LEAD Stainless steel is an alloy of IRON, CARBON and CHROMIUM (a transition metal) 30

31 Reactions of Metals Metal + Water Metal + Acid Metal + Oxygen Metal hydroxide + Hydrogen Salt + Water Metal oxide Reactivity Series By observing the reactions of metals we are able to build the Reactivity Series which shows how reactive metals are relative to each other. Extraction of Metals Some metals are uncombined. This means we find them pure in the ground; not joined to other elements. Examples of uncombined metals are silver and gold. Most metals are not like this. We find them combined with other elements in a compound known as an ORE. Most ores are metals joined to oxygen (metal oxide). To obtain a metal from a metal oxide we need to separate the metal atoms from the oxygen atoms. The more reactive the metal the tighter it is joined to the oxygen so the harder it is to produce! 31

32 Extraction of Metals Batteries In a battery electricity comes from a chemical reaction. Batteries need to be replaced when the chemicals inside are used up. Some batteries are rechargeable; e.g. the lead-acid battery. All batteries contain electrolytes. (Usually ammonium chloride paste) The purpose of the electrolyte is to complete the circuit. Mains Power v. Batteries Advantages Disadvantages Mains Power Cheaper Danger of being electrocuted Uses finite resources Batteries Safe to use Portable Expensive 32

33 The Electrochemical Series (ECS) Shows metals in order of their willingness to give up electrons. Shown on page 7 of the data booklet (Reactivity Series). Cells Electricity can be produced by connecting two different metals together (with an electrolyte eg ammonium chloride) to form a cell. Electricity can also be produced in a cell by connecting two different metals in solutions of their metal ions. Electrons flow from the metal higher in the ECS to the metal lower i.e. from magnesium to lead. The purpose of the ion bridge is to complete the circuit. The greater the distance between the metals in the electrochemical series, the higher the voltage produced. Displacement Reactions Reactions which occur when a metal higher up in the electrochemical series is added to a solution containing ions of a metal lower down in the series. E.g. Magnesium metal added to copper ions; Mg(s) + Cu 2+ SO 4 2- (aq) Mg 2+ SO 4 2- (aq) + Cu(s) The higher magnesium metal will form ions in solution The lower copper ions will form copper metal and come out of solution. I.e. they will be displaced. 33

34 The ECS can be used to predict whether or not a displacement reaction will occur. The metal being added must be higher than the ions in solution for displacement to occur. Corrosion Corrosion is a chemical reaction which involves the surface of a metal changing from an element to a compound (metal oxide). This is an example of oxidation. The corrosion of iron is known as rusting. Water and oxygen are required for corrosion to occur. Preventing Corrosion A surface barrier to air (oxygen) and water can provide physical protection against corrosion. E.g. Painting Greasing Coating with plastic Galvanising - metal objects are dipped into molten zinc Electroplating - silver, chromium and other metals can be deposited on the surface of a metal Tin-plating Gas Tests - Summary Gas Oxygen Hydrogen Carbon dioxide Test Relights a glowing splint Burns with a pop Turns limewater milky/cloudy 34

35 Electrical Conductivity Unit 7 Bonding, Structure & Properties Electric current is a flow of charged particles Testing conductivity 1. electrons flow through metals. 2. ions flow through solutions or melts. Testing a solid Testing a solution Metal elements and carbon (graphite) are electrical conductors. Non-metal elements are non-conductors of electricity. State at Room Temperature Covalent compounds do not conduct in any state as they only contain non-metals. Ionic compounds conduct when molten or in solution. Ionic compounds do not conduct when solid. Covalent compounds can be solids, liquids or gases. Ionic compounds are all solids. 35

36 Structure of Compounds Ionic solids exist as lattices of oppositely charged ions. E.g. Na + Cl - Some covalent solids are network structures E.g. SiO 2, diamond Ionic lattices and covalent networks have strong bonds and therefore high melting points. Other covalent compounds exist as discrete molecules. (see opposite) Discrete covalent molecules have weak forces of attraction and so these compounds have low melting points. Solubility Some covalent substances do not dissolve in water but will dissolve in other covalent solvents. Eg. Nail varnish being dissolved in propanone. Paint dissolving using turpentine. Ionic compounds generally dissolve in water. Eg. Sodium chloride & copper sulphate. Dissolving ionic compounds break up the ionic lattice. Solid ionic lattice Dissolved ionic compound (ions are free to move) 36

37 Electrolysis Electrolysis is used to break up a compound into its elements using electricity. Positive metal ions are attracted to the negative electrode. Negative non-metal ions are attracted to the positive electrode. Electrodes are made from carbon in the form of GRAPHITE as it conducts electricity. A d.c. (direct current) supply must be used if the products are to be identified. Covalent compounds cannot be electrolysed as they do not conduct electricity and they have no ions. Coloured Ions Some ions are coloured. Examples are shown in the table below. Ion copper nickel potassium chromate permanganate sulphate Colour blue green colourless yellow purple colourless 37

38 Migration of Ions In the experiment below; the positively charged blue copper ions move towards the negative electrode the negatively charged orange dichromate ions move towards the positive electrode. 38

39 Glossary Definition Acid Substance with a ph less than 7. Acid rain Source of pollution caused by sulphur dioxide SO 2 and nitrogen dioxide gas NO 2 Alkali Substance with a ph greater than 7. Alkali metals Very reactive group 1 elements. Atomic number Number of protons in the nucleus of an atom. Base Substance which will neutralise an acid. Blast furnace Catalyst Chemical reaction A tall oven used to extract iron from iron ore by burning it with carbon at high temperatures. A substance that increases the rate of a chemical reaction without being used up. A change in which new substances are made and cannot easily be reversed. Combustion Reaction which involves burning a substance in oxygen (O 2 ). Concentration The number of molecules of a substance in a given volume. Units are mol/l, mol l -1 or M Concentrated A solution that contains a large amount of solid dissolved. solution Corrosion Surface of a metal changing from a metal to a compound (metal oxide). Covalent bond A shared pair of electrons between two non-metal atoms. Crude oil A fossil fuel made up of a mixture of hydrocarbons. Diatomic Containing only 2 atoms. E.g. HCl, H 2, NO, CO, O 2 Dilute solution A solution containing a small mass of dissolved solute. Electrolysis Breaking a compound into its elements using electricity. Electron Negatively charged subatomic particle found in electron shells. Electron arrangement Element Endothermic Exothermic Fossil fuel Shows the number of electron shells and how many electrons they contain. A substance made up of only one type of atom. A reaction which releases heat to the surroundings ie gets cold. A reaction which releases heat. Fuels formed from fossils over millions of years at a high temperature and pressure; coal, oil and gas. 39

40 Fractional distillation Group Graphite Halogen Hydrocarbon Method of separating crude oil depending on boiling point. Vertical column of elements in the Periodic Table, with similar properties. Form of carbon with is an electrical conductor. Used as electrodes. Group 7 elements in the Periodic Table. A molecule containing only hydrogen and carbon. Insoluble Ion Ionic bond A substance than cannot be dissolved in water. A charged particle formed when an atom loses or gains electrons. A bond formed between a metal and a non-metal. Loam Metal hydroxide Metal oxide Metamorphic rock Mixture Mohs scale Molecule Noble gases Nucleus Neutral Period Proton Precipitation Reactive Reactivity Rusting Saturated solution Soluble Solute Solution Solvent A mixture of small pieces of rock or sand surrounded by decayed animal or plant remains. A compound containing metal, hydrogen and oxygen atoms which will neutralise an acid and has a ph greater than 7. A compound containing metal and oxygen atoms only. Rock formed when heat and pressure cause changes in existing igneous or sedimentary rocks over a long period of time. Two or more substances brought together but are not chemically joined. Eg. Air and crude oil. Scale used to measure the hardness of a mineral. A small group of atoms that are held together by covalent bonds. Group 8/0 elements in the Periodic Table. Unreactive elements with a full outer shell of electrons. The small, dense, postively charged centre of an atom, made up of protons and neutrons. Substance with a ph 7 which turns universal indicator green. Horizontal line going across the Periodic Table. Positive subatomic particle found within the nucleus. When two solutions react and a solid is formed. A substance that reacts quickly or easily. Eg. Alkali metals. How quickly or easily a substance will react. The specific name for the corrosion of iron. A solution in which no more solute can be dissolved. A solute which can dissolve in a liquid. A solid which is dissolved in the solvent. A liquid containing a dissolved solid. A liquid which can dissolved a solute. 40

41 Transition metals Universal Indicator Unreactive Valency Metals that are found in the middle of the periodic table. They do not have a group number given to them. Their valency is given in Roman numerals. Eg. Titanium. Solution used to test the ph of substances. Acids red, alkali blue and neutral green. A substance that reacts very slowly or does not react at all. Eg. Noble gases. How many bonds an atom can form. Eg. Group 4 elements have a valency (combining power) of 4 and therefore form 4 bonds. 41