Chapter 4 Quantities of Reactants & Products

1. Chapter 4 Quantities of Reactants & Products General Chemistry I Talia Ara

2. A. Chemical Equations (Reactions) • Understanding chemical reactions is one of the fundamental pursuits of chemistry. Chemical reaction: a process by which substances change into other substances by rearrangement, combination, or separation of atoms eg. A burning candle is an example of a combustion reaction.

3. A. Chemical Equations (Reactions) • A chemical equation is the symbolic representation of a chemical reaction in terms of chemical formulas. 4 Fe (s) + 3 O2 (g) 2 Fe2O3 (s) eg. The reaction of powdered Iron with oxygen to form iron oxide is represented by the equation above.

4. 1. Writing a Chemical Equation • a) Reactants & Products: In a chemical equation, the reactants (starting materials) are listed on the left & the products are listed on the right. • The reaction proceeds from left to right. ReactantsProducts (In words, the arrow means “yields”)

5. 1. Writing a Chemical Equation • b) Physical States: the physical states of the reactants and products in the reaction are indicated in parentheses after each chemical formula 2 Na (s) + Cl2 (g) 2 NaCl (s) (s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous solution

6. 1. Writing a Chemical Equation • c) Coefficients: the relative amounts of each substance in the reaction are indicated with stoichiometric coefficients 2 Na (s) + Cl2 (g) 2 NaCl (s) eg. Notice there are the same number of Na atoms on each side of the equation, and the same number of Cl atoms.

7. 2. Types of Chemical Reactions • There are far too many chemical reactions to memorize each and every one. • Many reactions fall into one of the simple reaction patterns shown below:

8. 2. Types of Chemical Reactions • a) Combination Reaction: reaction in which two or more substances combine to form one product • Oxygen and the halogens (Group 7A) frequently undergo combination reactions.

9. a) Combination Reaction Combination reaction between zinc and iodine: Zn (s) + I2 (s) ZnI2 (s)

10. 2. Types of Chemical Reactions • b) Decomposition Reaction: reaction in which one substance decomposes to form two or more products (the opposite of a combination reaction) • Some compounds that are stable under normal conditions will decompose upon heating.

11. b) Decomposition Reaction Decomposition of nitroglycerine: 4 C3H5(NO3) 3 (l) 12 CO2 (g) + 10 H2O (g) + 6 N2 (g) + O2 (g) • Nitroglycerine is very sensitive to heat, light and shock. • It can decompose violently.

12. 2. Types of Chemical Reactions • c) Displacement Reaction: reaction in which one element reacts with a compound to form a new compound and release a different element • The element released is said to have been displaced.

13. c) Displacement Reaction Displacing H with Na: The Reaction of Sodium Metal with Water

14. 2. Types of Chemical Reactions • d) Exchange Reaction: reaction in which there is an interchange of partners between two compounds • Exchange reactions commonly occur between ionic compounds dissolved in water (more in the next chapter).

15. d) Exchange Reaction Exchange Reaction Between Lead Nitrate and Potassium Chromate Pb(NO3)2 (aq) + K2CrO4 (aq) PbCrO4 (s) + 2 KNO3 (aq) lead nitrate potassium chromate lead chromate potassium nitrate

16. B. Balancing Chemical Equations • The Law of Conservation of Massdictates that the total number of atoms of each element on each side of a chemical equation must match. - An equation that meets this criterion is said to be a balanced equation.

17. B. Balancing Chemical Equations • For example: CaCO3 (s) CaO (s) + CO2 (g) Reactants:Products: 1 Ca atom 1 Ca atom 1 C atom 1 C atom 3 O atoms 3 O atoms

18. 2. Rules for Balancing Chemical Equations Step 1: Write an unbalanced equation containing the correct formulas of all the reactants and products.

19. 2. Rules for Balancing Chemical Equations Step 1: Write an unbalanced equation containing the correct formulas of all the reactants and products. eg. Ammonia from hydrogen and nitrogen H2 (g) + N2 (g) NH3 (g)

20. 2. Rules for Balancing Chemical Equations • Step 2: Use a whole-integer coefficient to balance the atoms of one of the elements. • Which element should you start with? • Start with the element that appears in the fewest formulas – preferably one that does not appear more than once on the same side of the equation. • Do not start with an element that is already balanced.

21. 2. Rules for Balancing Chemical Equations Step 2: Use a whole-integer coefficient to balance the atoms of one of the elements. H2 (g) + N2 (g) NH3 (g)

22. 2. Rules for Balancing Chemical Equations Step 2: Use a whole-integer coefficient to balance the atoms of one of the elements. H2 (g) + N2 (g) 2 NH3 (g) -Adding a 2 to NH3 balances the nitrogen atoms at two per side. - A coefficient multiplies all atoms in the corresponding formula – there are 2 N atoms and 6 H atoms on the right hand side of the equation.

23. 2. Rules for Balancing Chemical Equations Step 3: Balance the atoms of all the remaining elements. This may involve some trial and error in more complicated reactions. H2 (g) + N2 (g) 2 NH3 (g)

24. 2. Rules for Balancing Chemical Equations • Step 3: Balance the atoms of all the remaining elements. This may involve some trial and error in more complicated reactions. • 3 H2 (g) + N2 (g) 2 NH3 (g) • Adding a 3 to H2 brings the total number of • hydrogen atoms per side to six.

25. 2. Rules for Balancing Chemical Equations Step 4: Verify that the number of atoms of each element is balanced. 3 H2 (g) + N2 (g) 2 NH3 (g)

26. 2. Rules for Balancing Chemical Equations Step 4: Verify that the number of atoms of each element is balanced. 3 H2 (g) + N2 (g) 2 NH3 (g) (3 x 2)H + 2 N = 2 N + (2 x 3)H 6 H + 2 N = 2 N + 6 H Balanced!

27. Some Practice… S8 (s) + F2 (g) SF6 (g)

28. Some Practice… S8 (s) + F2 (g) SF6 (g) S8 (s) + F2 (g) 8 SF6 (g)

29. Some Practice… S8 (s) + F2 (g) SF6 (g) S8 (s) + F2 (g) 8 SF6 (g) S8 (s) + 24 F2 (g) 8 SF6 (g)

30. Some Practice… S8 (s) + F2 (g) SF6 (g) S8 (s) + F2 (g) 8 SF6 (g) S8 (s) + 24 F2 (g) 8 SF6 (g) 8 S + (24 x 2)F = 8 S + (8 x 6)F 8 S + 48 F = 8 S + 48 F Balanced!

31. Some Practice… Al(OH)3 (s) Al2O3 (s) + H2O (g)

32. Some Practice… Al(OH)3 (s) Al2O3 (s) + H2O (g) 2 Al(OH)3 (s) Al2O3 (s) + H2O (g)

33. Some Practice… Al(OH)3 (s) Al2O3 (s) + H2O (g) 2 Al(OH)3 (s) Al2O3 (s) + H2O (g) 2 Al(OH)3 (s) Al2O3 (s) + 3 H2O (g)

34. Some Practice… Al(OH)3 (s) Al2O3 (s) + H2O (g) 2 Al(OH)3 (s) Al2O3 (s) + H2O (g) 2 Al(OH)3 (s) Al2O3 (s) + 3 H2O (g) 2 Al + (2 x 3)O + (2 x 3)H = 2 Al + (3 + 3)O + (3 x 2) H 2 Al + 6 O + 6 H = 2 Al + 6 O + 6 H Balanced!

35. Some Practice… Fe (s) + Cl2 (g) FeCl3 (s)

36. Some Practice… Fe (s) + Cl2 (g) FeCl3 (s) Fe (s) + 3 Cl2 (g) 2 FeCl3 (s)

37. Some Practice… Fe (s) + Cl2 (g) FeCl3 (s) Fe (s) + 3 Cl2 (g) 2 FeCl3 (s) 2 Fe (s) + 3 Cl2 (g) 2 FeCl3 (s)

38. Some Practice… Fe (s) + Cl2 (g) FeCl3 (s) Fe (s) + 3 Cl2 (g) 2 FeCl3 (s) 2 Fe (s) + 3 Cl2 (g) 2 FeCl3 (s) 2 Fe + (3 x 2)Cl = 2 Fe + (2 x 3) Cl 2 Fe + 6 Cl = 2 Fe + 6 Cl Balanced!

39. Some Practice… C5H12 (g) + O2 (g) CO2 (g) + H2O (g)

40. Some Practice… C5H12 (g) + O2 (g) CO2 (g) + H2O (g) C5H12 (g) + O2 (g) 5 CO2 (g) + H2O (g)

41. Some Practice… C5H12 (g) + O2 (g) CO2 (g) + H2O (g) C5H12 (g) + O2 (g) 5 CO2 (g) + H2O (g) C5H12 (g) + O2 (g) 5 CO2 (g) + 6 H2O (g)

42. Some Practice… C5H12 (g) + O2 (g) CO2 (g) + H2O (g) C5H12 (g) + O2 (g) 5 CO2 (g) + H2O (g) C5H12 (g) + O2 (g) 5 CO2 (g) + 6 H2O (g) C5H12 (g) + 8 O2 (g) 5 CO2 (g) + 6 H2O (g)

43. Some Practice… C5H12 (g) + O2 (g) CO2 (g) + H2O (g) C5H12 (g) + O2 (g) 5 CO2 (g) + H2O (g) C5H12 (g) + O2 (g) 5 CO2 (g) + 6 H2O (g) C5H12 (g) + 8 O2 (g) 5 CO2 (g) + 6 H2O (g) 5 C + 12 H + (8 x 2) O = 5 C + (6 x 2) H + (10 + 6) O 5 C + 12 H + 16 O = 5 C + 12 H + 16 O Balanced!

44. Some Practice… Ni(OH)2 (aq) + HNO3 (aq) Ni(NO3)2 (aq) + H2O (l) Try this one at home…

45. C. Stoichiometry • Stoichiometry:the study of the quantitative relationships between amounts of reactants and products in a chemical reaction • Stoichiometric Coefficients: the multiplying numbers assigned to the species in a chemical equation in order to balance the equation CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)

46. 1. The Macro-Nano Connection • On the nanoscale, the stoichiometric coefficients in an equation represent the relative numbersof molecules or atoms in a reaction. • CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g) 1 molecule 2 molecules 1 molecule 2 molecules

47. 1. The Macro-Nano Connection • On the macroscale, the coefficients represent the molar ratios, the relative number of moles of reactants and products in a reaction. • These molar ratios can be used to calculate the molar amount of one compound from the molar amount of another. CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g) 2 mol O21 mol CO22 mol H2O 1 mol CH4 1 mol CH4 1 mol CO2

48. 1. The Macro-Nano Connection • By incorporating the molar masses of the compounds, the relative massesof the reactants and products in a balanced chemical equation can be calculated. • Notice the total mass of the reactants must equal the total mass of the products in a balanced equation.

49. a) Using Molar Ratios N2 (g) + 3 H2 (g) 2 NH3 (g) Using this balanced equation, determine the number of moles of NH3 that could be obtained from 3.6 mol H2.

50. a) Using Molar Ratios N2 (g) + 3 H2 (g) 2 NH3 (g) The coefficients in the balanced equation represent the molar ratios, so: