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C2 Additional Chemistry

Explore the concepts of ionic and covalent bonding, their properties, structures, and applications in various fields. Learn about isotopes, relative mass, and percentage yield.

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C2 Additional Chemistry

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  1. C2 Additional Chemistry

  2. Ionic Bonding

  3. Atoms want to have a full outer shellof electrons to become stable just like the noble gases H Non-Metals Metals don’t have enough have too many electrons

  4. Ionic Bonding X X F Li X Metal atoms lose electrons to become positive ions. Non-metals atoms gain electrons to become negative ions. Oppositely charged ions have a strong electrostatic attraction towards each other. -1 X +1 F Li X X

  5. After ionic bonding we sometimes get differently charged ions...but compounds are always neutral! So to find the formula combine ions and make sure that the charges cancel each other out. E.g. Na2O, MgCl2, Al(OH)3, Al2O3

  6. Ionic Properties Oppositely charged ions are attracted to each other to form a giant ionic lattice. A lot of heat energy is needed to break the strong bonds so ionic compounds have high melting points. Ionic compounds can also conduct electricity when they are molten or in a solution because ions are free to move.

  7. Metallic Bonding

  8. Before Bonding After Bonding Bonding & Structure Metals lose their outer electrons to become positively charged ions These ions form layers to give a giant metallic lattice These are surrounded by a sea of delocalised electrons The delocalised electrons are free to move

  9. Properties of Metals - - Delocalised e-s are free to move - + + + - - • High melting points due to strong attraction between positive ions and electrons • Conduct electricity and heat because the electrons are free to move + - - - - + + + + + -

  10. Pure metals: • One type of atom only • Regular layers • Layers can slide easily • Malleable (soft) Alloys: • Mixture of metals • Distorted layers • Cannot slide easily • Much harder Pure Metal Alloy

  11. Covalent Bonding

  12. Covalent Bonding • Non-metal atoms only • Atoms share electrons • A molecule is formed

  13. Simple molecules Simple covalent molecules consist of a set number of atoms. e.g. H2O, CO2, O2 They have weak intermolecular bonds between the molecules. Little energy is needed to break these bonds so they have low boiling points. O O O C C O Weak bonds

  14. Giant Structures Giant covalent structures consist of infinite atoms. E.g. Sand is made from silicon dioxide Strong covalent bonds  veryhigh melting points Carbon atoms exists as 3 different giant structures Graphite Diamond Silicon dioxide Fullerene

  15. Diamond: • Each carbon atom forms 4 covalent bonds • Strong bonds hold atoms in a rigid, pyramid structure • Extremely hard – used for cutting tools

  16. Graphite: • Each carbon forms 3 covalent bonds • Flat, layered structure • Weak intermolecular forces between layers • Layers can slide over each other - soft & slippery • Delocalised electrons are free to move - conducts electricity

  17. Properties of Polymers

  18. = = = = Polymers are long molecules that are formed when lots of small monomers are joined together by covalent bonds. The bonds/forces between the polymer chains determines their properties - - - - - - - - - - - - - - - - - - - -

  19. Types of poly(ethene)

  20. Thermosofteningvs Thermosetting Thermosofteningplastics become soft & melt when heated. They have weak forces between the polymer chains. Forces are easily broken allowing the chains to slide over each other. Thermosetting plastics do not become soft when heated. They have strong covalent cross links between the polymer chains. These bonds do not break, keeping the chains together.

  21. Nanoscience

  22. Nanoscience is the study of tiny nanoparticles Nanoparticles contain just a few hundred atoms and are less than 100nm in size These have an extremely high surface area to volume ratio and have special properties and uses: Cosmetics and suncream Electronics Catalysts Gold nanoparticles used in cancer therapy Nanotubes used in electronic products

  23. Relative Mass

  24. Mass of atoms Li 7 N N 3 The top number on the periodic table tells us the mass of an atom It tells us the number of protons +neutrons. (electrons have no mass) P P N P N neutrons = mass – atomic number (top – bottom number)

  25. Isotopes Cl Cl 35 37 17 17 Protons = 17 Neutrons = 18 Isotopes are forms of the same element which have a different relative atomic mass. They have the same number of protons but a different number of neutrons. Protons = 17 Neutrons = 20

  26. Relative Atomic Mass Cu 63.5 Cl O 15.99 35.5 17 8 29 The relative atomic mass, Ar, of an element is the average mass of its atoms compared to the isotope 12C. Ar is often rounded to the nearest whole number.

  27. Relative Formula Mass Mg H C O 24 1 12 16 12 1 6 8 Relative formula mass (Mr) = Sum of the relative atomic masses in a compound • MgO = 24 + 16 = 40 • H2O = 1 + 1 + 16 = 18 • CO2 = 12 + 16 + 16 = 44

  28. Percentage Mass & Yield

  29. Percentage Mass C O 12 16 relative atomic mass of element in question Small number after element in question 8 6 % mass = Ar x number of atoms x 100 Mr of compound Relative molecular mass of whole compound E.g. Find the % mass of O in CO2 % mass = 16 x 2 x 100 12+16+16 = 72.7%

  30. Percentage Yield We do not always get as much product from a reaction as we expect. This may be because some of the product is left on the apparatus or spilt. % yield = Actual mass x 100 Expected mass E.g. 6g of ammonia was produced even though 24g was expected to be produced % mass = (6÷24) x 100 = 25% Reversible reactions give very low yields!!

  31. Empirical Formulae

  32. The empirical formula tells you the number of each type of atom in a compound. to find... • Show formula as a ratio • Divide the masses, or percentages, by the relative atomic masses (Ar) • Simplify ratio by dividing all answers by smallest number • Use numbers to write empirical formula

  33. S O 32 16 16 8 Examples • A compound contains 64g of sulphur and 64g of oxygen. What is it’s empirical formula? • S : O • 64÷32 : 64÷16 2 : 4 • 2÷2 : 4÷2 1 : 2 • SO2

  34. Examples • A compound contains 40% Ca, 12% C and 48% O. What is it’s empirical formula? Relative atomic masses,Ar : C=12, O=16, Ca=40 • Ca : C : O • 40÷40 : 12÷12 : 48÷16 1 : 1 : 3 • Already simplified! • CaCO3

  35. Mole Calculations

  36. Moles tells us how many particles are in a substance. In 1 mole there are 6x1023 particles. A mole is basically just a really big number, sometimes known as Avogadro’s number Na 23 Mass (g) 11 Ar or Mr Moles E.g. Find the mass in grams of 2 moles of Na Mass = moles x Ar = 2 x 23 = 46g

  37. Mole ratios can be used together with the moles triangle to calculate the number of moles & masses of all reactants & products in a reaction. The number in front of each element or compound in a balanced equation gives us the mole ratio. Remember to lay your answer out in this way... 2Mg + O2 2MgO Mole ratio 2 : 1 : 2 Mass (g) Some of these will be in the Q Relative mass Calculate but ignore mole ratio (ignore big numbers) Moles Calculate using triangle or ratio

  38. Mass (g) Ar or Mr 1. What mass of water is produced when 20g of hydrogen reacts with oxygen? Relative atomic masses (Ar): H=1, O=16 2H2+ O2  2H2O Mole ratio 2 : 1 : 2 Mass (g) 20 180 Relative mass 2 32 18 Moles 105 10 Moles x Found using mole ratio

  39. Mass (g) Ar or Mr 2. What mass of water is needed to make 49g of sulphuric acid? Relative atomic masses (Ar): H=1, O=16, S=32 SO3+ H2O  H2SO4 Mole ratio 1 : 1 : 1 Mass (g) 9 49 Relative mass 80 18 98 Moles 0.50.5 0.5 Moles x Found using mole ratio

  40. Summary of Calculations Mass (g) Relative atomic mass (Ar): Top number (periodic table) Relative formula mass (Mr): Sum of atomic masses Ar or Mr Moles Percentage mass = Ar x number of atoms x 100 Mr of compound % yield= Actual mass x 100 Possible mass Empirical Formula: 1) Show formula as a ratio 2) Divide percentages by the Ar 3) Simplify ratio 4) Write empirical formula

  41. Analysing Substances

  42. Paper Chromatography Most Soluble Least Soluble Paper chromatography can be used to separate mixtures. The number of spots shows the number of substances. The height of the spots can be used to identify the substance

  43. Using Paper Chromatography 1 2 3 4 mix 1 2 3 4 mix The mix consists of dyes 1 and 4

  44. Analysing Mixtures Using GCMS GC-MS machines can be used to analyse mixtures of compounds. They are used to test for drugs in an athlete’s urine. Gas chromatography (GC) separates the mixture Mass spectrometry (MS) measures the relative molecular mass of the compounds. C9H13NO Mr = 151

  45. Sample is injected into a gas stream • Gas stream carries sample through a long thin column • Some compounds take longer to pass through the column Gas chromatography (GC)

  46. Molecular ion peak shows the relative molecular mass Mass spectrometry (MS)

  47. Pros and Cons of GCMS

  48. Speeding up reactions

  49. Collision theory Particles need to collide with enough energy for a reaction to take place. This is called the activation energy. A B A B C

  50. Temperature Increase temperature Particles have more kinetic energy Particles collide more often and with more energy More collisions above the activation energy Rate of reaction increases

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