2. When a substance dissolves in water, it forms an aqueous solution (aq). The solution can be acidic, neutral or alkaline. The indicators are used to tell if a solution is acidic, alkaline or neutral. They become different colours with different solutions.

3. alkaline solutionacidic solution litmusbluered phenolphthaleinpinkcolourless methyl orangeyellowred

4. pH is a measure of how acidic or how alkaline a solution is. The pH scale goes from 0 to 14, with 0 being very strongly acidic, and 14 being very strongly alkaline. A pH of 7 is neutral.

5. pHcharacteristics 0 – 3strongly acidic 4 – 6weakly acidic 7neutral 8 – 10weakly alkaline 11 - 14strongly alkaline

6. Universal indicator It can show the strengths of the acids and alkalis because it has more colours. Each colour is linked to a number called the pH scale.

7. An acid is any substance which produces H + ions. An alkali is any substance which produces OH - ions. OH - ions are called hydroxide ions.

8. A substance which will neutralise an acid, but does not dissolve in water, is called a base. For example, copper(II) oxide, iron(II) oxide and zinc carbonate are bases, they do not dissolve in water.

9. ChlorideCl - NitrateNO 3 - SulfateSO 4 2- CarbonateCO 3 2- HydroxideOH -

10. Reactants 1. dilute hydrochloric, nitric and sulfuric acid 2. metals, metal oxides and metal carbonates Metals: groups 1, 2, 3

11. acid+ metal→ metal salt + hydrogen acid+ metal oxide → metal salt + water acid+ metal carbonate → metal salt + water + carbon dioxide Except nitric acid reacting with metals. Equations.

12. General rules for predicting the solubility of salts in water 1. All common sodium, potassium and ammonium salts are soluble. 2. All nitrates are soluble.

13. 3. Common chlorides are soluble, except silver chloride. 4. Common sulfates are soluble, except those of barium and calcium. 5. Common carbonates are insoluble, except those of sodium, potassium and ammonium.

14. How to prepare soluble salts from acids acid (aq) + alkali (aq)→ salt (aq) + water (l) acid (aq) + base (s)→ salt (aq) + water (l) acid (aq) + carbonate (s) → salt (aq) + water (l) + carbon dioxide (g) acid (aq) + metal (s)→ salt (aq) + hydrogen (g) Equations.

15. Precipitation reactions The process of making a solid come from a solution is called precipitation. The solid itself is called a precipitate. An insoluble salt can be made by reacting the appropriate soluble salt with an acid or alkali or another salt.

16. calcium chloride + sodium carbonate → calcium carbonate + sodium chloride CaCl 2(aq) + Na 2 CO 3(aq) → CaCO 3(s) + 2NaCl (aq) The ionic equation is: Ca 2+ (aq) + CO 3 2- (aq) → CaCO 3(s) A precipitate can be separated from the solution by filtration. The precipitate can then be left somewhere warm to dry.

17. Titration A titration is used to make soluble salts or to calculate concentration. How much acid is needed to neutralise an alkali? In a titration an acid can be added drop by drop to an alkaline solution. The pH of the alkaline solution will decrease as more acid is added.

18. When the pH goes down to 7, the solution will be neutral. If more acid is added, the pH will continue to go down and the solution becomes acidic. The point where the solution becomes neutral (called the end point) can be found using an indicator.

19. Using an indicator The exact amount of acid needed to neutralise an alkali can be found by titration. A titration uses a burette, pipette and a conical flask. In the example below, an acid and an alkali react to make sodium chloride.

20. hydrochloric acid + sodium hydroxide → sodium chloride + water HCl (aq) + NaOH (aq) → NaCl (aq) + H 2 O (l) The burette is filled with hydrochloric acid. A known quantity of alkali is released from a pipette into the conical flask. The tap on the burette is turned open to allow the acid to be added drop by drop into the alkali.

21. The alkali contains an indicator (e.g. phenolphthalein). When enough acid has been added to neutralise the alkali the indicator changes from pink to colourless. Universal indicator is not usually used for a titration because it changes gradually giving different colours for a different pH.

22. Methyl orange or phenolphthalein are used because they give a sudden change in colour at neutralisation which makes it easier to see the end point of the titration.

26. Endothermic and Exothermic A chemical reaction always has a change in energy. In a reaction going from reactants to products, either 1) Heat energy is given out - called exothermic. 2) Heat energy is taken in - called endothermic. Melting and boiling are endothermic, freezing and condensing are exothermic.

27. Breaking of bonds requires energy, you have to put heat in - it is endothermic. This is why melting and boiling are endothermic. Making of bonds gives out energy - it is exothermic. This is why freezing and condensing are exothermic.

28. In a chemical reaction you need to put energy in to break bonds in the reactants, you get energy out when new bonds are formed to make the products.

29. If you get out more energy than you have to put in, then overall the reaction is exothermic. If you have to put in more energy than you get out, then the reaction is endothermic.

30. Calculating energy transfers The amount of energy transferred from the reaction to the water in the calorimeter can be calculated if you know: the mass of water heated the temperature rise

31. For a given amount of water heated up, the greater the temperature rise, the greater the amount of heat energy transferred to the water. For example, twice as much energy is transferred to the water to achieve a temperature increase of 20 ºC compared with 10 ºC.

33. Measuring heat transfers is called calorimetry. Simple calorimetry experiments: combustion, displacement, dissolving and neutralisation.

34. Combustion: spirit burner Displacement: metal + acid Dissolving: see worksheet Neutralisation: strong acid and strong alkali

35. The change in energy during a reaction is given the symbol ∆H. (represents molar enthalpy change for exothermic and endothermic reactions) OR

36. ΔH is the difference between the energy needed to break the bonds in the reactants, and the energy given out when new bonds are made in the products.

37. For an exothermic reaction, ∆H is negative. For an endothermic reaction, ∆H is positive. The change in energy can be plotted against the progress of a reaction, as the reactants turn into products.

39. Going from reactants to the top of the curve, you are going up the energy scale. Energy 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.

40. Going from the top of the curve to the products, you are going down the energy scale. Energy is given out as bonds form in the products.

41. The reactants are higher up the energy scale than are the products. The amount of energy you need to put in is less than the amount of energy you get out. This is a typical exothermic reaction.

45. Collision theory Collision theory says that a chemical reaction can only occur between particles when they collide. Particles may be atoms, ions or molecules. There is a minimum amount of energy which colliding particles need in order to react with each other.

46. If the colliding particles have less than this minimum energy then they just bounce off each other and no reaction occurs. This minimum energy is called the activation energy.

47. The faster the particles are going, the more energy they have. Fast moving particles are more likely to react when they collide. You can make particles move more quickly by heating them up (raising temperature).

48. Changing the rate of a reaction There are 5 ways to increase the rate of a chemical reaction. They are all understood in terms of collision theory. The rate of a chemical reaction may be increased by 1) Raising the temperature. 2) Increasing the concentration (in solution). 3) Increasing the pressure (in gases).

49. 4) Increasing the surface area of a solid. 5) Use a catalyst. The opposite of 1, 2, 3 and 4 will decrease the rate of a reaction. A catalyst can make a reaction go faster or slower. In practice a catalyst is mainly used to make a reaction go faster.

50. Measuring the rate The rate of a reaction may be measured by following: the loss of a reactant, or the formation of a product. Rate of reaction = amount of reactant used ÷ time Rate of reaction = amount of product formed ÷ time

51. They are 1. The reaction between calcium carbonate and dilute hydrochloric acid. 2. The reaction between sodium thiosulfate solution and hydrochloric acid. 3. The decomposition of hydrogen peroxide solution. http://www.absorblearning.com/chemistry/demo/units/LR15 01.html

52. The reaction between calcium carbonate and dilute HCl 2HCl (aq) + CaCO 3(s) → CaCl 2(aq) + CO 2(g) + H 2 O (l) The rate of this reaction can be measured by following the rate at which carbon dioxide is formed.

53. This can be done by conducting the reaction in an open flask on an electric balance. As the carbon dioxide escapes to the air, the mass of the flask will decrease. You can take a reading from the balance every 30 seconds, then plot a graph of loss of mass against time. The gradient of the plot shows the rate of the reaction.

54. The reaction between sodium thiosulfate solution and dilute hydrochloric acid 2HCl (aq) + Na 2 S 2 O 3(aq) → 2NaCl (aq) + SO 2(g) + S (s) + H 2 O (l)

55. The solid sulfur (S (s) ) formed in this reaction makes the colourless solution go cloudy. The reaction is usually carried out in a flask placed on a piece of white paper which has a black cross on it. At the beginning of the reaction, the cross can be seen easily.

56. As the flask becomes more and more cloudy the cross gets harder to see. You can measure the time from the start of the reaction until the cross can no longer be seen. This is a way of measuring the rate of formation of sulfur.

57. The decomposition of hydrogen peroxide solution 2H 2 O 2(aq) →O 2(g) +2H 2 O (l)

58. The reaction is carried out in a closed flask which has a gas syringe connected to the top of it. The reaction is started by adding a catalyst to the hydrogen peroxide. The catalyst is manganese (IV) oxide.

59. The volume of oxygen in the syringe increases as the reaction proceeds. The volume of oxygen can be noted every 30 seconds, and a graph of volume against time can be plotted. The gradient of the plot shows how fast the reaction is going.

60. Raising the temperature Raising the temperature makes the particles move faster. This means that the particles collide more frequently with each other and the rate of the reaction increases. Also, the faster the particles are travelling, the greater is the proportion of them which will have the required activation energy for the reaction to occur.

61. Increasing the concentration (in solution) Increasing the concentration of a substance in solution means that there will be more particles per dm 3 of that substance. The more particles that there are in the same volume the closer to each other the particles will be. This means that the particles collide more frequently with each other and the rate of the reaction increases.

62. Increasing the concentration of sodium thiosulfate means that the solid sulfur will be produced more quickly and there will be less time before the cross can no longer be seen. When making different concentrations of sodium thiosulfate take care to use the same total volume of sodium thiosulfate plus hydrochloric acid for the comparison to be fair.

64. Increasing the pressure (in gases) Increasing the pressure of a reaction where the reactant is a gas is similar to increasing the concentration of a reactant in a solution. The gas particles will be closer together when the pressure increases.

65. This means that the particles collide more frequently with each other and the rate of the reaction increases. If the reaction is reversible, then increasing the pressure will favour the side of the reaction which has the smaller volume.

66. Increasing the surface area of a solid A solid in a solution can only react when particles collide with the surface. The bigger the area of the solid surface, the more particles can collide with it per second, and the faster the reaction rate is. You can increase the surface area of a solid by breaking it up into smaller pieces. A powder has the largest surface area and will have the fastest reaction rate. This is why catalysts are often used as powders.

68. Rates of reaction presentation

69. Temperature and rate http://www.chemactive.com/captivate/grade_11/rate_temper ature.htm

70. Surface area and rate http://www.chemactive.com/captivate/grade_11/rate_ar ea.htm

71. Concentration and rate http://www.chemactive.com/captivate/grade_11/rate_concen tration.htm

72. Catalysts and rate http://www.chemactive.com/captivate/grade_11/rate_catalys t.htm

74. When a reaction is reversible, it means that it can go both forwards and backwards. The forward reaction is the one we want, in which reactants are converted into products. The backward reaction is where the products become changed back into the original reactants. Reactions in both directions occur at the same time.

75. Heating copper(II) sulfate crystals Anhydrous copper(II) sulfate is white. Anhydrous copper(II) sulfate is copper(II) sulfate which is completely dry. When water is added to it, it turns into hydrated blue crystals. This is used as a test for water.

76. If blue (hydrated) copper(II) sulfate crystals are heated, they lose their water, turn white and become anhydrous copper(II) sulfate crystals again. The reaction is reversible.

77. Heating ammonium chloride If you heat ammonium chloride, the white crystals disappear from the bottom of the test tube and reappear further up. Heating ammonium chloride splits it into the colourless ammonium and hydrogen chloride. These gases recombine further up the tube, where it is cooler. The reaction is reverse when the conditions are changed from hot to cool. http://www.absorblearning.com/media/item.action?quick=1tp

78. If the forward reaction is exothermic, the backward reaction is endothermic. The amount of heat transferred in the forward reaction is the same as the amount of heat transferred in the backward reaction. In a closed system, after awhile an equilibrium mixture is reached, where a certain proportion of the mixture exists as reactants, and the rest exists as products.

79. When equilibrium has been reached, it does not mean that the reactions have stopped. It means that the forward reaction is making products at the same rate that the backward reaction is making reactants. It is said to be a dynamic equilibrium.

80. For a reversible reaction, Le Chatelier's Principle states that "The equilibrium position will respond to oppose a change in the reaction conditions". What this means in practice is 1) If you remove a product, the equilibrium mixture changes to make more product. It tries to get back to the composition it had before the product was removed.

81. 2. Heat may be treated as a reactant (for an endothermic reaction) or as a product (for an exothermic reaction). If you remove heat from an exothermic reaction, the equilibrium will change to produce more product.

82. This will not only produce more heat, but also produce more of the chemical product that you want in the equilibrium mixture. If you add heat to an exothermic reaction, the reverse will happen, and you will get less product in the equilibrium mixture.

83. 3. For a reversible reaction involving gases, increasing the pressure will change the equilibrium to make more of the side of the reaction which has the smaller volume. Decreasing the pressure will change the equilibrium to make more of the side of the reaction which has the larger volume.