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C2 revision powerpoint

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Summary of each specification point for Edexcel C2.

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  • what is this? you said that metal becomes positively charged and become anions in slide 6, then you said they become cations in slide 8? Im confused bruuhhh
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  • Mendeleev created the periodic table in 1869
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C2 revision powerpoint

  1. 1. C2 Topic 1 and 2
  2. 2. C2.1 Mendeleev • Mendeleev created the periodic table in 1871. • He arranged elements in order of their properties. • -groups = down. • -period= across • Metals are on the left hand side of the periodic table. He arranged them by the following properties: • Atomic mass • Density • Melting Point • Formula of the oxide Patterns emerged down the groups and across the periods which confirmed his predictions.
  3. 3. C2.2 Structure of the atom • Atoms are made up of PROTONS, NEUTRONS and ELECTRONS. • They are found in the positions shown on the diagram. • Atom is the smallest of particles. Electron shells Nucleus containing protons and neutrons Particle Relative Mass Relative Charge Proton 1 +1 Neutron 1 0 Electron 1/1840 -1 Electrons fill shells from the middle in the order of 2, 8, 8, 18 (how many elements are in each period) The group number corresponds to the number of electrons in the outer shell.
  4. 4. C2.3 The Modern Periodic Table• Top number – MASS NUMBER The number of protons and neutrons • Bottom number – ATOMIC NUMBER The number of protons (also the same as electrons) Al 27 13 Protons = 13 Electrons = 13 Neutrons = 27-13 = 14 Protons and neutrons are packed together tightly in the nucleus (high density) Electrons are spread out in shells (low density) Isotope - Elements with the same number of protons and electrons but different number of neutrons.
  5. 5. C2.4 Electron shells • Electron located in the shell. • Electronic configuration = how the electrons are arranged in an atom.
  6. 6. C2.5 Ionic Bonding • Ionic bonds form between METALS and NON-METALS. • Ionic bonding involves the transfer of ELECTRONS. • Metallic Ions are POSITIVELY charged (ANIONS). • Non-metallic elements are NEGATIVELY charged (CATIONS). Loose electron +
  7. 7. Common Ions Charged Particles. To become positively charged an atom must lose electrons. To become negatively charged and atom must gain electrons.
  8. 8. C2.5 Ionic compounds- non metal and a metal. • Conduct electricity when MOLTEN (melted) and in an AQUEOUS SOLUTION (dissolved in water) • DO NOT conduct electricity as a SOLID • Have high MELTING and BOILING points • Usually SOLID at ROOM TEMPERATURE Ion = an atom with a positive or negative charge. Cations = metal atoms lose electrons to form positively charged ions called cations. Anions = Non-metal atoms gain electrons to form negatively ions called anions. Transferring electrons = ionic bonds. Ionic compounds have a lattice structure, with a regular arrangement of ions, held together by electrostatic forces between oppositely charged ions.
  9. 9. C2.6 Ionic compounds • Ionic compounds contain ions. Below are some of the common ones. Compound ions = contain more than one element. Names of ionic compounds = If the compound contains oxygen it ends with an – ate…..eg….. Iron sulfate = Fe SO4 If no oxygen then they end in an –ide…….Iron Sulfide….FeS
  10. 10. C2.7 Properties of ionic Compounds • Solid = does not conduct electricity • Liquid = conduct electricity…WHY….ions are free to move around and they are charged. • Melting points = when a substance changes from solid to liquid. • Boiling point = changes from a liquid to a gas. • Ionic substances = high boiling points and high melting points. • Ionic bonds = strong and lots of them to break so needs lots of energy.
  11. 11. C2.8 SolubilitySoluble = substance dissolves. Insoluble = does not dissolve Salt = Acid + alkali Table explains which salts are soluble and which are not.
  12. 12. 2.10 Precipitations. silver nitrate (soluble) + sodium chloride (soluble) → silver chloride (insoluble) + sodium nitrate (soluble) AgNO3 (aq) + NaCl (aq)→ AgCl (s) + NaNO3 (aq) Precipitation reaction - An insoluble solid formed when two soluble substances are mixed. Barium sulfate is used as a ‘barium meal’ to x-ray patients because it is 1. Opaque to x-rays 2. Safe to use 3. Insoluble so doesn’t enter the blood
  13. 13. C2.10 Precipitates • Precipitates are SALTS that are formed in chemical reactions that DO NOT DISSOLVE in the solvent used in the reaction.
  14. 14. C2.11 Ion Tests • Ion tests for metals (ANIONS) are usually done by FLAME TESTS. • Each ION produces a certain flame colour. Testing for CATIONS is done through chemical testing. 1. Chloride – add nitric acid and silver nitrate, if a white precipitate forms then chloride ions are present. 2. Sulphate – add hydrochloric acid and barium chloride, if a white precipitate forms then sulphate ions are present. 3. Carbonate – add an acid and if the gas produced turns limewater MILKY the carbonate ions were present. (carbonate release carbon dioxide)
  15. 15. C2 Topic 3 & 4
  16. 16. C2.12 Covalent Compounds Learn these covalent molecules. Be able to draw them using dot and cross diagrams. Covalent bond is when a pair of electrons are shared between two atoms forming a molecule. Methane CH4 Water H2O Hydrogen H2 Carbon Dioxide CO2 Hydrogen Chloride HCl Oxygen O2
  17. 17. C2.13 Properties of Covalent Substances Simple Covalent Substances • Low melting point • Low boiling point • Poor conductor of electricity Why? • Because there are weak forces between molecules. Giant Covalent Substances • High melting point • High boiling point Why? • Because there are strong forces between molecules.
  18. 18. 2.13 Diamond vs. Graphite Diamond • Used to make cutting tools. • Very hard with strong bonds between molecules. • Does not conduct electricity because there are no free electrons. Graphite • Used to make electrodes and lubricant. • There are strong forces between molecules within the layers – but there are weak forces between the layers. • There is one free electron for each carbon atom meaning it can conduct electricity. Both examples of giant covalent molecules. Both made from CARBON.
  19. 19. C2.15 Miscible or Immiscible? Key definitions Miscible – Liquids that completely mix to form a solution. Immiscible – Liquids that do not form a solution but separate into two layers e.g. oil and water. Separating two immiscible liquids Use a separating funnel. 1. Open the tap and let the lower liquid flow out and collect in a beaker. 2. Close the tap before the upper liquid starts to run out. 3. Change the beaker and then let the other liquid flow out. Separating two miscible liquids Use fractional distillation to separate liquid air to form nitrogen and oxygen. Exploits the fact they have different boiling points. 1. Cool air to -200oC to liquefy it. 2. Warm up to -185oC to evaporate the nitrogen and keep the oxygen liquid.
  20. 20. C2.16 Chromatography Used to separate mixtures of coloured compounds including inks, dyes and colouring agents in food. 1. A spot of the mixture is placed near the bottom of a piece of chromatography paper and the paper is then placed upright in a suitable solvent, e.g. water. 2. As the solvent soaks up the paper, it carries the mixtures with it. 3. Different components of the mixture will move at different rates. 4. This separates the mixture out. Rf values Different chromatograms and the separated components of the mixtures can be identified by calculating the Rf value using the equation: The Rf value of a particular compound is always the same - if the chromatography has been carried out in the same way. This allows industry to use chromatography to identify compounds in mixtures. Rf = distance moved by the compound ÷ distance moved by the solvent
  21. 21. C2.17 Chemical Classification Type of bonding Melting Point Boiling Point Solubility in Water Conductivity Ionic High High Most Dissolve Conduct electricity when molten or in solution, not when solid Simple molecular covalent low low Some dissolve in water Do not conduct electricity Giant molecular covalent High High Insoluble Do not conduct except graphite Metallic High except mercury High except mercury Insoluble Good conductors
  22. 22. C2.14 Classifying Substances Substance Melting Point Boiling Point Electrical Conductivity Solubility in water Type of Bonding Sodium Chloride High High Yes in solution or molten Soluble IONIC Magnesium Sulfate High High Yes in solution or molten Soluble IONIC Hexane Low Low No Insoluble SIMPLE COVALENT Paraffin Low Low No Insoluble SIMPLE COVALENT Silicon dioxide High High No Insoluble GIANT COVALENT Copper sulphate High High Yes in solution or molten Soluble IONIC Sucrose High High Soluble GIANT COVALENT Could you classify these substances based on their properties? Know these examples!
  23. 23. C2.18 Metallic Bonding and Transition Metals The structure of metals – a regular arrangement of positive ions surrounded by a sea of delocalised electrons. Metals are malleable because ions can slide over each other if a large enough force is applied. The electrons hold the ions together meaning the metal spreads out rather than breaks. Metals conduct electricity because the delocalised electrons are free to move. If a voltage is applied, the electrons all move in the same direction and this is called a current. Most metals are transition metals and they typically have high melting points and produce coloured compounds.
  24. 24. C2.19 Alkali Metals Alkali metals are soft with relatively low melting points. Lithium, Sodium and Potassium react with water to form hydroxides (which are alkaline) and hydrogen gas. Alkali metals get more reactive as you move down the groups because the atoms lower down have more electron shells. This means the outer electron is further away from the nucleus, not held in place as strongly, and easier to lose – making it more reactive. Lithium + water  Lithium Hydroxide + Hydrogen Sodium + water  Sodium Hydroxide + Hydrogen Potassium + water  Potassium Hydroxide + Hydrogen
  25. 25. C2.20 Halogens Halogen Colour State at room temperature Fluorine Pale Yellow Gas Chlorine Green/yellow Gas Bromine Brown Liquid Iodine Grey Solid Metal + Halogen  Metal Halide Potassium + Bromine  Potassium Bromide 2K(s) + Br2(l)  2KBr2(S) Calcium + Chlorine  Calcium Chloride 2Ca(s) + Cl2(l)  2KCl2(S) They all follow the same pattern. Learn them.
  26. 26. C2.21 Displacement of the halogens • Displacement reactions – a more reactive element will displace (take the place of) a less reactive element. Potassium Bromide + Chlorine  Potassium Chloride Potassium Chloride + Bromine  No reaction Chlorine is more reactive than bromine so will displace the bromine. Sodium Iodide + Chlorine  Potassium Chloride Sodium Chloride + Iodine  No reaction Chlorine is more reactive than iodine so will displace the iodine.
  27. 27. C2.22 More halogen reactions Halogens react with hydrogen to make hydrogen halides which will dissolve in water to make acidic solutions. Hydrogen + Fluorine  Hydrogen Fluoride H2(g) + F2(g)  2HF(g) Hydrogen + Chlorine  Hydrogen Chloride H2(g) + Cl2(g)  2HCl(g) Salt Sodium fluoride Sodium chloride Sodium bromide Sodium iodide Halogen Fluorine Reaction Reaction Reaction Chlorine No reaction Reaction Reaction Bromine No reaction No reaction Reaction Iodine No reaction No reaction No reaction Displacement reactions can be used to work out the relative reactivity of halogens.
  28. 28. C2.23 Noble gases Noble gases are chemically inert (unreactive) because they have a full outer shell of electrons. Discovery of Noble Gases 1. Chemists noticed the density of nitrogen made in a reaction was different from that of nitrogen in the air. 2. Chemists made a hypothesis about the composition of the air. 3. Chemists performed tests to show the presence of the Noble gases in the air. Uses of Noble Gases Reason Xenon and Argon inside filament lamps To stop the hot filament reacting with oxygen and burning away Argon and Helium used in welding To form a blanket over the hot metal to stop it reacting with oxygen in the air Argon used in fire extinguishing systems It is non flammable and can be used to fill a computer server room (for example) if a fire breaks out. Helium used to fill balloons and air ships It has low density Neon in fluorescent lamps Produces a red light when electric current is passed through it under low pressure
  29. 29. C2 Topic 5 Chemical Reactions
  30. 30. C2.24 and C2.25 Temperature changes Keywords • Endothermic – reaction that takes heat energy in, decreasing the temperature of the reaction mixture and its surroundings • Exothermic – reaction that releases heat energy, increasing the temperature of the reaction mixture and its surroundings Facts: • During a chemical reaction there is usually a transfer of energy between the reactant and the surroundings. Endothermic • Not many reactions are endothermic • E.g. Sodium hydrogen carbonate + Hydrochloric acid Sodium Chloride + Water + Carbon dioxide • Other endothermic reactions include: – Photosynthesis – Dissolving ammonium nitrate • Takes in heat energy Exothermic • Gives out heat • Most reactions are exothermic • Temperature increase • All combustion reactions are exothermic • E.g. Methane + Oxygen • Explosions – release a lot of heat and gases very quickly Making/breaking bonds • Energy needed to break bond • Energy released when bonds made
  31. 31. C2.26, 2.27 and 2.28 Rates of Reactions and Collision Theory Keywords • Rate of reaction – The speed at which a reaction takes place • Concentration – A measure of how much solute is dissolved in a fixed volume of solvent. • Surface area– The total area of all the surfaces of an object or substance Factors affecting Rate 1. Concentration • More concentrated = more particles • More particles = more collisions = faster reaction 2. Temperature • Higher temperature = faster reaction e.g. And egg cooks faster in boiling water than warm water • Particles have more energy = move faster – More effective collisions (collide with more energy) – Collide more frequently 3. Surface area (SA) • Solid broken up into smaller pieces = larger SA • Greater surface area = faster reaction • More surface area = more particles on the surface therefore more frequent collisions • A = Smaller SA (block) • B = Larger SA (powder) Examples • Fast reactions = Burning, explosions • Slow reaction = Rusting, apple browning
  32. 32. Catalysts • Many chemical processes use catalysts to increase rate of production of products • Catalysts help to lower the temperature and pressure needed = less energy needed = saves money Keywords • Catalyst – A substance that speeds up the rate of a reaction without being used up in the reaction • Catalytic converter – Device fitted to car exhausts with a thin layer of transition metal catalyst on a honeycomb structure giving a large surface area.. The catalyst speeds up the reaction to combine Carbon monoxide and unburned petrol into carbon dioxide and water. C2.29 Catalysts Catalytic converters • Reduce pollutants in exhaust gases • Combine Carbon monoxide (CO) with oxygen • Carbon dioxide and water released instead • Contain transition metals (platinum, rhodium or palladium) • Expensive metals so small amounts used • Catalyst spread over a honeycomb structure = large Surface Area • Works faster with the hot gases from the engine heat it up
  33. 33. C2 Topic 6 Quantitative Chemistry
  34. 34. C2.30 and 2.31 Relative Mass RAM • Compares mass to that of carbon 12 • E.g. Hydrogen is 12X lighter that carbon RFM • Sum of all masses in the formula • E.g. EF • Simplest whole number ratio of atoms/ions • E.g. H2O there are double the number of hydrogens to oxygen Keywords • Relative Atomic mass (RAM) – The mean mass of an atom relative to the mass of an atom of carbon-12, which has a mass of 12. • Empirical formula (EF) – The simplest whole number ration of atoms of each element in a compound • Molecular formula (MF) - The actual of atoms of each element that combine to make a molecule of a compound. • Relative formula mass (RFM)– The sum of the relative atomic masses of all the atoms in a formula
  35. 35. B2.32 Percentage composition • No atoms are lost or made in a chemical reaction. • They are rearranged into new substances • You can use relative mass and balanced equations to calculate the mass of a reactant or product. Percentage by mass of an element in a compound A M 100 Number of atoms of element = X X
  36. 36. C2.33 Yields Products need to be make as cheaply as possible. Chemists need to make sure the reaction creates as much product as possible. Theoretical Yield 1.Maximum calculated amount of a product that could be formed from a given amount of reactants. 2.Can be calculated from the balanced equation 3.Assumes all reactant are turned into products and all products harvested Actual Yield 1.The actual amount of product obtained from a chemical reaction. Percentage Yield 1.The actual yield divided by the theoretical yield as a percentage Yield usually less than expected – 3 reasons: 1. Reaction may be incomplete 2. Some product is lost Keywords Yield – the amount of useful product obtained from a reaction.
  37. 37. C2.34 Waste and Profit Cost Effective Process • All factories must consider environmental impacts including disposal of waste. • Companies want the most profit • Reactions need to be • High in percentage yield • All products are useful so there are no waste products (or uses found for waste products – econimically viable) • Reactions are fast (quick to make lots of product) Disposal of Waste Products • Industry make many useful substances (cement, pesticides, plastics) • Many of the reactions make waste products:  Not useful  Must be disposed of following strict laws  Expensive to dispose  Can cause environmental problems  Can cause social problems (house prices drop or unpleasant smells) Keywords • By-products – Any product formed in a reaction in addition to the required product • Waste products – By-products that have no uses

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