1. Quantitative Relationships in Chemical Change CHEMISTRY 20

2. BALANCING CHEMICAL REACTIONS Every chemical reaction involves the rearrangement of atoms into different combinations. However, during these reactions, the total number of atoms of each type of element is the same after the reaction as it was before the reaction. Chemical reactions have to be properly balanced in order to clearly obey the Law of Conservation of Matter. Every chemical reaction must first be written so that each reactant and product has the correct chemical formula and state of matter. Coefficients are then used in order to balance the various atoms. All coefficients must be in the simplest whole number ratio possible.

3. BALANCING CHEMICAL REACTIONS Hints for balancing equations: Write out an unbalanced equation, making sure that each reactant and product formula is written correctly. Use coefficients to balance any atoms that are not already balanced. If a polyatomic ion is remaining intact, it may be easiest to balance it as a group. Always balance any elements last. Make sure that the coefficients have the simplest whole number ratio possible.

4. EXAMPLES Re-write the following word equations as balanced chemical equations. sodium + chlorine → sodium chloride Na (s) + Cl (s) → NaCl (s) aluminium chlorate → aluminium + chlorine + oxygen

5. EXAMPLES butane(C 4 H 10(g) ) + oxygen → carbon dioxide + water vapour scandium + copper(II) sulfate → copper + scandium sulfate hydrochloric acid + barium hydroxide → water + barium chloride

6. FORMATION REACTIONS A formation reaction is a reaction in which two or more elements react together to form a compound. X + Y → XY N 2(g) + 3 H 2(g) → 2 NH 3(g)

7. Write a balanced chemical equation for the formation of glucose (C 6 H 12 O 6(s) ). Write a balanced chemical equation for the formation of ammonium benzoate. EXAMPLE

8. DECOMPOSITION REACTIONS A decomposition reaction is a reaction in which a compound reacts and breaks down into its component elements. XY → X + Y 2 H 2 O (l) → 2 H 2(g) + O 2(g)

9. Write a balanced chemical equation for the decomposition of diphosphorous heptaoxide. Write a balanced chemical equation for the decomposition of sodium sulfate. EXAMPLE

10. COMBUSTION REACTIONS A complete combustion reaction is a reaction in which a hydrocarbon burns in a plentiful supply of oxygen and the only products are carbon dioxide and water vapour. CH 4(g) + 2 O 2(g) → CO 2(g) + 2 H 2 O (g)

11. Write a balanced chemical equation for the complete combustion of hexane (C 6 H 14(l) ). Write a balanced chemical equation for the complete combustion of methanol (CH 3 OH (l) ). EXAMPLE

12. HOMEWORK Balancing Chemical Reactions practice questions page 5 in your notes

13. SINGLE REPLACEMENT REACTIONS

14. Write a balanced chemical equation for the reaction between zinc and hydrochloric acid. Write a balanced chemical equation for the reaction between aluminium chloride and bromine. EXAMPLE

15. DOUBLE REPLACEMENT REACTIONS

16. Write a balanced chemical equation for the reaction between aluminium nitrate and ammonium phosphate. Write a balanced chemical equation for the reaction between hydrosulfuric acid and potassium hydroxide. EXAMPLE

17. CALCULATING AMOUNTS

18. EXAMPLE Calculate the amount of potassium carbonate in a 33.5 g sample.

19. EXAMPLE Calculate the amount of zinc nitrate that must be dissolved to prepare 750 mL of a solution that has a concentration of 15.8 mmol/L.

20. EXAMPLE Calculate the amount of helium in a 8.00 L balloon if it has a pressure of 103 kPa at 22.0 o C.

21. HOMEWORK Balancing Chemical Reactions practice questions Page 8-9 in your notes

22. NET IONIC EQUATIONS Chemistry 20

23. NET IONIC EQUATIONS In replacement reactions, many of the reactants and products exist as dissociated ions. Some of these dissociated ions remain unchanged in any way throughout the reaction.

25. DEMO Potassium iodide + lead (II) nitrate https://www.youtube.com/watch?v=DITY2rXYU-I

26. DEMO

28. EXAMPLE Write a complete balanced equation, ionic equation, and net ionic equation for the reaction between barium hydroxide and hydrochloric acid.

29. EXAMPLE Write a complete balanced equation, ionic equation, and net ionic equation for the reaction between aluminium and copper(II) nitrate.

30. QUALITATIVE VS QUANTITATIVE ANALYSIS Qualitative analysis involves determining by experiment whether a certain substance is present in a sample. Quantitative analysis involves determining how much of a certain substance is present in a sample. For aqueous solutions, typical qualitative analytic techniques include observing solution colour, flame tests and precipitation reactions. When certain ions are dissolved in water, they give the solution a distinct color.

31. QUALITATIVE VS QUANTITATIVE ANALYSIS The more concentrated the ion is in the solution, the more evident this characteristic colour will appear. Many metal ions produce a distinct colour of flame when they are heated. One way to test for the presence of metal ions in solution is to heat a drop of the solution in a hot flame and observe the colour. This is called a flame test. The different colours that fireworks can have are due to the explosive ignition of different metals and metal salts.

32. HOMEWORK Practice Questions on page 11 and 14 in your notes

33. SELECTIVE PRECIPITATION Chemistry 20

34. SELECTIVE PRECIPITATION Precipitation reactions can determine whether an ion is present in solution or not. A precipitation reaction is another term for a double- replacement reaction in solution that produces a solid product (the precipitate ). Chemists add dissolved substances to unknown solutions and observe whether a precipitate forms. https://www.youtube.com/watch?v=PaA5LOawjpE

35. At each stage, the colour of the solution is observed and the resulting precipitate is removed. Flame tests can be used to identify the precipitates.

36. EXAMPLE You are given an unidentified solution and are told that it may or may not contain sulfide ions. How could you confirm or deny the presence of S 2- (aq) in this solution?

37. EXAMPLE You are given an unidentified solution and are told that it may or may not contain zinc ions. How could you confirm or deny the presence of Zn 2+ (aq) in this solution?

38. If a solution contains more than one dissolved ion, it is essential to design the technique carefully so that only one precipitate is formed.

39. EXAMPLE You are given an unidentified solution and are told that it may or may not contain acetate ions and/or sulfate ions. How could you confirm or deny the presence of CH 3 COO - (aq) and/or SO 4 2- (aq) in this solution?

40. EXAMPLE You are given an unidentified solution and are told that it may or may not contain lead(II) ions and/or aluminium ions. How could you confirm or deny the presence of Pb 2+ (aq) and/or Al 3+ (aq) in this solution?

41. Flame test demo Homework: Practice Questions pg 17/18 Remember you have a quiz on Wednesday, up to and including today!

42. STOICHIOMETRY Chemistry 20

43. STOICHIOMETRY

45. Because of the large numbers of molecules involved in any chemical reaction, it is more convenient to compare the amount (number of moles) of the reactants and products. For every 2 mol of NO 2(g) that are produced, 2 mol of NO (g) and 1 mol of O 2(g) have been consumed. In a reaction, the actual amounts involved may vary but this ratio will always be observed.

46. EXAMPLE

47. What amount of zinc will be produced by the decomposition of 0.40 mol of zinc phosphate?

48. EXAMPLE During the formation of benzoic acid, 0.366 mol of carbon is consumed. What amount of hydrogen will have been consumed during this reaction?

49. HOMEWORK Practice problems in your notes pg 21-22

50. GRAVIMETRIC STOICHIOMETRY Chemistry 20

51. GRAVIMETRIC STOICHIOMETRY

52. STOICHIOMETRIC PROCESS 1.Write a balanced chemical equation. 2.Using the information given, calculate the amount of the given substance ( n given ) by the following equation: 3.Calculate the amount of the required substance ( n required ) using the mole ratio from the balanced equation and the n given you calculated in step 2. 4.Calculate the mass of the required substance using the n required you calculated in step 3 by the following equation:

53. EXAMPLE What mass of oxygen must be available in order to burn 120 g of ethane (C 2 H 6(g) )?

54. EXAMPLE Solid lithium hydroxide reacts with carbon dioxide gas to produce lithium carbonate and water vapour. What mass of lithium carbonate will be produced when 5.00 g of lithium hydroxide is used up?

55. EXAMPLE Solutions of sodium bromide and lead(II) acetate are mixed together. The precipitate is filtered, dried, and found to have a mass of 2.17 g. What minimum mass of lead(II) acetate was dissolved in the original solution?

56. HOMEWORK Practice problems in your notes pg 25-26

57. SOLUTION STOICHIOMETRY Chemistry 20

58. SOLUTION STOICHIOMETRY

59. EXAMPLE What volume of 0.214 mol/L sodium hydroxide would be required to completely neutralize 500 mL of 0.0104 mol/L hydrochloric acid?

60. EXAMPLE A 100 mL portion of hydrochloric acid is able to react with 5.00 g of zinc. What is the concentration of the hydrochloric acid solution?

61. EXAMPLE A student mixes 225 mL of 0.078 mol/L cobalt(II) nitrate with an excess volume of sodium hydroxide. Predict the mass of precipitate that should be made.

62. HOMEWORK Practice problems in your notes pg 28-29

63. GAS STOICHIOMETRY Chemistry 20

64. GAS STOICHIOMETRY

65. EXAMPLE A solution of hydrochloric acid is able to react completely with 5.00 g of zinc. What volume of hydrogen gas will be produced at 21.0 o C and 99.6 kPa?

66. EXAMPLE A sample of cyclopentane (C 5 H 10(l) ) is burned completely. During the reaction, 125 L of oxygen at SATP is consumed. What mass of cyclopentane has been burned?

67. HOMEWORK Practice problems in your notes page 31-32

68. LIMITING AND EXCESS REACTANTS Chemistry 20

69. STOICHIOMETRIC AMOUNTS The coefficients of a balanced chemical equation are often called the stoichiometric coefficients because they are used in stoichiometric calculations. If the reactants are present in the amounts that correspond exactly to the mole ratios, they are said to be present in stoichiometric amounts. When the reactants are in stoichiometric amounts, then absolutely no trace of any of the reactants will be left at the end of the reaction.

70. In most reactions, one of the reactants may run out before the others. There will usually be one or more of the reactants left over without getting a chance to completely react. In these cases, the amount of product that results from a chemical reaction is limited by the reactant that is used up or completely consumed first. The reactant that is completely used up in the reaction is called the limiting reactant. It is also known as the limiting reagent. Any reactant(s) that are left over are called the excess reactant. The limiting reactant does not need to be the reactant present in fewer moles. Rather, it is the reactant that will form fewer moles of product(s).

71. LIMITING REACTANTS

72. EXAMPLE A 1.25 g piece of magnesium is placed into 80.0 mL of 0.113 mol/L hydrochloric acid. Which reactant is the limiting reactant? What amount of the excess reactant will remain unreacted at the completion of the reaction?

73. EXAMPLE Calculate the mass of precipitate that should be produced when 200 mL of 0.118 mol/L iron(II) sulfate is mixed with 175 mL of 0.204 mol/L ammonium phosphate.

74. HOMEWORK Practice questions pg 34 in your notes

75. PREDICTED AND EXPERIMENTAL YIELD Chemistry 20

76. PREDICTED AND EXPERIMENTAL YIELD Stoichiometry calculations can be used to predict the maximum quantity of product expected from a reaction. This quantity is known as the predicted yield (which is also known as the theoretical yield ). The predicted yield is calculated on the assumption that all the limiting reactant reacts to make product on the ratio described by the balanced equation. The quantity of product actually obtained by a reaction is called the experimental yield (which is also known as the actual yield ). In most reactions, the experimental yield will not match exactly with the predicted yield. Usually, it is a lower value than the predicted yield.

77. FACTORS THAT LIMIT EXPERIMENTAL YIELD

78. Slow Reaction: If a reaction is slow and not enough time has been allowed for the reaction to reach completion, the quantity of products measured will be less than predicted.

79. FACTORS THAT LIMIT EXPERIMENTAL YIELD Collection and Transfer Methods: If a precipitate is collected by filtration, some of it may remain dissolved in the filtrate. When a precipitate is rinsed to remove traces of the reactants, some of the precipitate may dissolve in the rinsing solvent. Mechanical losses are the small amount of product that are lost when they remain stuck to glassware or filter paper as they are transferred in the lab.

80. FACTORS THAT LIMIT EXPERIMENTAL YIELD Reactant Purity: Many chemicals used in the laboratory that are reactant-grade may be close to 100% pure, but there may be trace amounts of contaminants.

81. FACTORS THAT LIMIT EXPERIMENTAL YIELD Reactions That Do Not Proceed To Completion: Many reactions reach a point where the reaction appears to stop, although less than 100% of the reactants have been converted into products. These reactions have reached equilibrium and the products are reacting to form the reactants at the same rate as the reactants are reacting to form the products. For example, under most conditions, only a small percentage of hydrogen and iodine molecules have reacted to form hydrogen iodide at any one time: H 2(g) + I 2(g)  2 HI (g)

82. CALCULATING PERCENT YIELD

84. EXAMPLE A student mixes together two solutions in the lab and forms a precipitate. The following data is recorded: Solution A: 85.0 mL of 0.172 mol/L potassium phosphate Solution B: 120.0 mL of 0.144 mol/L calcium nitrate Mass of filter paper: 1.14 g Mass of filter paper + dried precipitate: 2.85 g Calculate the percent yield for this reaction.

85. PRACTICE

86. ACID-BASE TITRATION Chemistry 20

87. ACID-BASE TITRATION In a titration, the concentration of one solution is determined by quantitatively observing its reaction with a standard solution (ie. a solution of known concentration). The observations can be used to standardize the solution (ie. determine its unknown concentration). The predicted yield is calculated on the assumption that all the limiting reactant reacts to make product on the ratio described by the balanced equation.

88. The solution that is placed into the burette is known as the titrant. By measuring the initial burette volume (prior to beginning the titration) and the final burette volume (at the completion of the titration), the volume of titrant required to complete the reaction can be determined. The solution that is placed into the Erlenmeyer flask is known as the aliquot. The volume of the aliquot is pre-determined before the reaction begins - a volumetric pipette is used to measure the aliquot. Either the aliquot or the titrant can be the standard solution. ACID-BASE TITRATION SET-UP

89. ACID-BASE TITRATION The stage of the titration at which the reaction is complete is called the equivalence point. At this point, stoichiometrically equivalent amounts of each reactant have been consumed. In acid-base titrations, an acid-base indicator is added to the aliquot to provide visual evidence of the end of the reaction. A dramatic colour change of the indicator identifies when the reaction is complete. The point at which the indicator changes colour is called the endpoint.

90. EQUIVALENCE POINT In an acid-base titration, an acid titrant is added to a base aliquot, or vice versa. For monoprotic acids and bases, the point at which equal moles of reactant acid and base combine is called the equivalence point of the titration.

91. EQUIVALENCE POINT For example, the titration of sodium hydroxide with hydrochloric acid. HCl (aq) + NaOH (aq) H 2 O (l) + NaCl (aq) The mole ratio is 1:1. Therefore, the equivalence point occurs when an equal amount of HCl (aq) has been added to the NaOH (aq).

92. In every reaction between a strong monoprotic acid and a strong base, the equivalence point has a pH of 7 because all hydronium ions from the acid have been neutralized by an equal amount of hydroxide ions from the base. Acid-base titrations are performed with repeated trials until at least 3 concordant results are obtained. A concordant result means the titrant volumes required to reach the equivalence point are within a range of 0.2 mL.

93. Most neutralization reactions involve colourless solutions with no obvious visible evidence that a reaction is taking place. An acid-base indicator is a substance that changes colour over a given pH range. Usually, indicators are weak monoprotic acids. The molecular and ionized forms of the indicator have different colours.

94. For example, bromothymol blue is a commonly used indicator for titrations. It is yellow between pH 0 and pH 6. It turns blue between pH 6 and pH 7.6. This indicator is commonly used for titrations between a strong monoprotic acid and a strong base.

95. TITRATION DEMONSTRATION

96. PH CURVES Chemistry 20

97. When a strong base titrant is reacted with a strong acid aliquot, the characteristic shape that should be predicted is:

98. When a strong acid titrant is reacted with a strong base aliquot, the characteristic shape that should be predicted is:

99. CHOOSING AN INDICATOR Titrations are usually performed with indicators because they are cheaper than pH meters and have easy to recognize colour changes at the equivalence point. The titration can only be done accurately if a suitable indicator is chosen. The endpoint pH of the indicator must be within the steep rise or drop in the titration curve. Ideally, the endpoint of the indicator would occur right at the equivalence point of the reaction.

100. Whenever a titration is between a strong monoprotic acid and a strong monoprotic base, the equivalence point will be observed at a pH of 7.00. Therefore, the most appropriate choice for indicator for this type of titration is either bromothymol blue or phenol red. CHOOSING AN INDICATOR