Chemistry revision!

1. The early periodic table The modern periodic table Group 1 – alkali metals Group 7 – halogens Transition elements Strong and weak acids and alkalis Titrations Titration calculations How ideas about acids and alkalis developed Water and solubility Solubility curves Hard water Removing hardness Water treatment Comparing the energy produced by fuels Energy changes in reactions Calculations using bond energies Test for positive ions Tests for negative ions Testing for organic substances Instrumental analysis 1 Instrumental analysis 2 Chemistry revision!

2. In the 19th Century some chemist dudes had a fight… • In the 1800s people were on a roll discovering elements. Scientists didn’t like how disorganised they were so they tried to organise them. The problem was they didn’t know a whole lot about them! • Not much was known about atoms • Each element had several names • A load of elements hadn’t been discovered yet!

3. John Dalton • Dalton put the elements in order of mass, measured by doing different chemical reactions. • Problem was, that didn’t tell anyone a lot about the elements and when more elements were discovered the list had to be changed!

4. John Newlands • Newlands introduced a bit more order; the law of octaves (8) • He based this on the fact that every eighth element seemed to have similar properties. • Problem is, he got a bit full of himself and tried to make all the elements fit in the octaves even if they didn’t fit the pattern, assuming all had been discovered in spite of the fact new ones were being discovered all the time. He even put 2 elements in one spot to make them all fit. • All the other scientists laughed at him .

5. Alexandre-Emile Beguyer de Chancourtois • This French dude copied Newlands idea of octaves, but put them in a much clearer diagram. • Unfortunately when his work was published the diagram was missed out! • Without the diagram everyone got confused so they ignored him 

6. Dmitri Mendeleev • This Russian dude cracked it. • By the time he ordered the elements 50 had been discovered. He put them in order of atomic masses and then arranged them so a periodic pattern in their physical and chemical properties could be seen. • The clever part was he left GAPS for elements that had not been discovered yet when the pattern wasn’t followed. He could predict the properties of these elements thanks to his table. • A few years later elements were discovered that fit Mendeleev’s predictions. • All the other scientists thought he was brilliant 

7. 1. The Early Periodic Table Dalton – elements in order of mass Newlands – law of octaves, organised into similar properties Mendeleev – periodic table used today. In order of atomic masses, and structured according to chemical and physical properties. He left gaps which were filled when new elements were discovered

8. 2. The modern Periodic Table • Elements are arranged in order of their atomic (proton) number, and in groups of similar properties. • Groups have the same number of electrons in their outer shell.

9. 3. Alkali Metals • VERY reactive, stored in oil (fizz and burn H2 with water) • Reactivity increases down the group • Low density (1st 3 float on water!) • Soft (can cut with a knife) • Low melting and boiling points for metals • Melting points decrease down the group • React with non-metals, losing their outer electron Lithium + water → lithium hydroxide + hydrogen 2Li (s) + 2H2O (l) → 2LiOH (aq) + H2(g) Sodium + chlorine → sodium chloride 2Na (s) + Cl2(g)→ 2NaCl (s)

10. 4. Halogens • Poisonous non-metals • Coloured vapours • Low melting and boiling points • Poor conductors of heat and electricity • Fluorine – yellow gas • Chlorine – green gas • Bromine – orange/brown liquid • Iodine – gray solid – violet vapour • Go around in pairs (F2, Cl2…) • Less reactive as you go down the group • H2(g) + F2(g)→ 2HF(g) • Displacement: Cl2 + 2KBr → 2KCl + Br2

11. 5. Transition Elements • Held together by metallic bonding • Good conductors of heat and electricity because of delocalised electrons • Hard, tough and strong • Malleable • High melting points (apart from mercury) • Much less reactive than alkali metals (corrode slowly) • Combining them makes useful alloys • Make coloured compounds

12. Task: • Page 222-3 • Summary questions 1 - 6

13. Strong and weak acids/alkalis • Acids form H+ ions when we add them to water HCl(g)→ H+(aq) + Cl-(aq) • The H+ (hydrogen ion) is the acidic part. • Acids are PROTON DONORS. A proton is an H+ • Form OH- ions when we add them to water NaOH(s) → Na+(aq) + OH-(aq) • OH- (hydroxide ion) is the alkaline part. • Alkalis are PROTON ACCEPTORS. • An acid or base is STRONG if it completely ionises in water. • An acid or base is WEAK if it only partially ionises in water.

14. 7. Titrations Indicators are used to show the endpoints of neutralisation reactions: • Strong acid + strong alkali - Universal indicator or any other • Weak acid + strong alkali - Phenolphthalein • Strong acid + weak alkali - Methyl orange In a titration we need to use the correct indicator We also use pipettes to measure out fixed volume of solution and burettes to measure the volume of solution added.

15. 8. Titration Calculations • On the bottles you find the name of the acid and alkali and a number followed by a M. • M means the number of moles in 1000cm3 (or 1dm3) of solution. ‘Molarity’ or ‘molar concentration’ mean the same thing. • n = M x V (dm3) for liquids

16. 9. How ideas about acids and alkalis developed • Liebig defined an acid as a compound that contained hydrogen which could react with a metal to produce hydrogen gas. • Arrhenius defined an acid as a substance that produces hydrogen ions (H+) in water, and bases as a substance that produces hydroxide ions (OH-) when dissolved in water. • As Arrhenius’ definition only worked in aqueous solution another definition had to be developed. • Brønsted and Lowry defined an acid as a proton donor and a base as a proton acceptor

17. Task Page 232-3 Summary Questions 1 – 3 Exam style questions 1 - 2

18. 10. Water and solubility • Soluble - substances that can dissolve. • Insoluble – substances that cannot dissolve. • Solution - the mixture formed when a substance dissolves. • Solute - the substance that dissolves. • Solvent - the liquid in the solution. A saturated solution is one in which no more solute will dissolve at that temperature. When a hot saturated solution cools some of the solute will separate from the solution.

19. 11. Solubility Curves • The solubility of most solid solutes increases as the temperature increases • The solubility of gases decreases as temperature rises. • The solubility of gases increases as pressure increases

20. 12. Hard water • Soft water readily forms lather with soap. Hard water reacts with soap to form scum and so more soap is needed to form a lather. • Hard water contains dissolved compounds, usually of calcium or magnesium. The compounds are dissolved when water comes into contact with rocks. • Using hard water can increase costs because more soap is needed. • Salts in the water react with the soap to produce stearates (scum) • When hard water is heated it can produce scale (calcium carbonate) that reduces the efficiency of heating systems and kettles. • Hard water has some health benefits because calcium compounds are good for health. • Calcium is good for strong teeth and bones. • There is evidence that drinking hard water reduces the chances of heart disease.

21. 13. Removing hardness • Hardness is from dissolved calcium and magnesium ions. The ions come from rocks which the water has filtered through. • The ions can be removed using washing soda (sodium carbonate), which precipitates the ions. • An ion exchange column can be used. They contain sodium ions which can be exchanged with the calcium or magnesium ions.

22. 14. Water Treatment • Water from boreholes is usually pretty clean, it’s been filtered by the surrounding rocks. Usually you just use chlorine to kill of germs. • Water from rivers or reservoirs needs more treatment. Treatment involves chemical processes, like adding aluminium sulphate and lime, and physical processes, like filtration. Water that has been treated is not pure. It still contains substances dissolved in it. Pure water is produced by distilling it, boiling and condensing the steam produced, or deionising it by using an ion exchange column.

23. Task: • Page 246-7 • Summary questions 1 - 3

24. 15. Comparing the energy produced by fuels • The relative amounts of energy released when substances burn can be measured by simple calorimetry, e.g. by heating water in a glass or metal container. This method can be used to compare the amount of energy produced by fuels and foods. • Energy is normally measured in joules (J). Some dietary information is given in calories, which are equal to 4.2 joules. Different foods produce different amounts of energy. Foods with higher proportions of carbohydrates, fats and oils produce relatively large amounts of energy.

25. 16. Energy changes in reactions • Energy (heat) is being put in to break bonds in the reactants. • At the top of the curve, the bonds in the reactants have been broken.The amount of energy put in to break these bonds is called the activation energy. • The activation energy is the minimum amount of energy needed for the reaction to occur. A catalyst may work by lowering the activation energy for a reaction. • Energy (heat) is given out as bonds form in the products. The difference in energy levels between the reactants and the products is given the symbol DH (pronounced 'delta H').This is the amount of heat given out (or taken in) during the reaction.For an exothermic reaction, DH is negative.For an endothermic reaction, DH is positive.

26. 17. Calculations using bond energies H2 + Br2→ 2HBr H – H Br – Br H – Br H - Br • H-H = 436 kJ/mol, Br-Br = 193jK/mol, H-Br = 366kJ/mol • Reactants: 436 + 193 = 629 kJ/mol • Product: 366 x 2 = 732 kJ/mol • 629 – 732 = -103 kJ/mol -ve means exothermic

27. Task: • Page 256-7 • Summary questions 1 - 2

28. 18. Tests for positive ions Add sodium hydroxide (NaOH) • Cu (II) – light blue ppt • Iron(II) – dirty green ppt • Iron(III) – red/brown ppt • Al, Ca, Mg – white ppt Add more NaOH to Al/Ca/Mg: • Al – white ppt dissolves again • Ca, Mg – white ppt does not dissolve • Add NaOH to ammonium ions (NH4+) • Makes ammonia! • Warm soln and use damp red litmus on gas – turns blue

29. Carbonates (CO32-): - Add dilute acid – makes CO2. - CO2 turns limewater cloudy/milky Copper carbonate turns green to black when heated Zinc carbonate turns white to yellow when heated (and kept hot!) Nitrates: Add NaOH and warm it. If no ammonia detected add Al and test for ammonia again (damp red litmus blue) Sulphates: Add hydrochloric acid and barium chloride. Makes white ppt (barium sulphate) Halides: Add dilute nitric acid and silver nitrate. Cl ions – white ppt Br ions – cream ppt I ions – yellow ppt 19. Tests for negative ions

30. 20. Testing for organic substances • Contain carbon • Burn or char on heating • You can detect C=C double bond with bromine water (orange to colourless) • COMBUSTION ANALYSIS allows you to determine the empirical formula of an organic substance.

31. 21. Instrumental Analysis Technology is awesome: • Highly accurate • Quicker • Enable small quantities to be analysed Technology has not so awesome: • Usually expensive • Takes special training to use • Gives results that can often be interpreted only by comparison with already available known specimens

32. 22. Instrumental analysis 2 • AAS – atomic absorption spectroscopy, measures concentrations of metals in liquids • Mass spectrometry – compares masses of different atoms using magnets! • UV – visible spectroscopy • NMR (nuclear magnetic resonance) spectroscopy Chromatography: • Gas-liquid – separates compounds easily vaporised • Gel permeation – separates according to size of molecules • Ion-exchange – separates according to charge • High performance liquid – separates compounds in solution