A Study of Chemical Reactions

1. A Study of Chemical Reactions Equations, Mole Conversions, & Stoichiometry

2. Types of Reactions • Many chemical reactions have defining characteristics which allow them to be classified as to type.

3. Types of Chemical Reactions • The five types of chemical reactions in this unit are: • Combination/Synthesis • Decomposition/Analysis • Single Replacement/Displacement • Double Replacement/Metathesis • Combustion

4. Combination Reactions • Two or more substances combine to form one substance. • The general form is A + X AX • Example: • Magnesium + oxygen  magnesium oxide • 2Mg + O2 2MgO

5. Magnesium + Oxygen

6. Combination Reactions • Combination reactions may also be called composition or synthesis reactions. • Some types of combination reactions: • Combination of elements • K + Cl2 • One product will be formed

7. Combination Reactions • K + Cl2 • Write the ions: K+ Cl- • Balance the charges: KCl • Balance the equation: 2K + Cl2  2KCl

8. Combination Reactions • Some types of combination reactions: • Oxide + water  • Nonmetal oxide + water  acid • SO2 + H2O  H2SO3 • Metal oxide + water  base • BaO + H2O  Ba(OH)2

9. Combination Reactions • Some types of combination reactions: • Metal oxides + nonmetal oxides • Na2O + CO2 Na2CO3 • CaO + SO2  CaSO3

10. Decomposition Reactions • One substance reacts to form two or more substances. • The general form is AX  A + X • Example: • Water can be decomposed by electrolysis. • 2H2O  2H2 + O2

11. Electrolysis of Water

12. Decomposition Reactions • Types of Decomposition Reactions: • Decomposition of carbonates • When heated, some carbonates break down to form an oxide and carbon dioxide. • CaCO3 CaO + CO2 • H2CO3  H2O + CO2

13. Decomposition Reactions • Types of decomposition reactions: • Some metal hydroxides decompose into oxides and water when heated. • Ca(OH)2 CaO + H2O Note that this is the reverse of a similar combination reaction.

14. Decomposition Reactions • Types of decomposition reactions: • Metal chlorates decompose into chlorides and oxygen when heated. • 2KClO3 2KCl + 3O2 • Zn(ClO3)2  ZnCl2 + 3O2 • Some of these reactions are used in explosives.

15. Decomposition Reactions • Some substances can easily decompose: • Ammonium hydroxide is actually ammonia gas dissolved in water. • NH4OH  NH3 + H2O • Some acids decompose into water and an oxide. • H2SO3 H2O + SO2

16. Decomposition Reactions • Some decomposition reactions are difficult to predict. • The decomposition of nitrogen triiodide, NI3, is an example of an interesting decomposition reaction.

17. Nitrogen triiodide

18. Single Replacement Reactions • Cationic: A metal will replace a metal ion in a compound. • The general form is A + BX  AX + B • Anionic: A nonmetal will replace a nonmetal ion in a compound. • The general form is Y + BX  BY + X

19. Single Replacement Reactions • Examples: • Ni + AgNO3 • Nickel replaces the metallic ion Ag+. • The silver becomes free silver and the nickel becomes the nickel(II) ion. • Ni + AgNO3 Ag + Ni(NO3)2 • Balance the equation: • Ni + 2AgNO3  2Ag + Ni(NO3)

20. Activity Series

21. Single Replacement Reactions • Not all single replacement reactions that can be written actually happen. • The metal must be more active than the metal ion. • Aluminum is more active than iron in Al + Fe2O3 in the following reaction:

22. Thermite Reaction

23. Thermite Reaction • Al + Fe2O3 • Aluminum will replace iron(III) • Iron(III) becomes Fe and aluminum metal becomes Al3+. • 2Al + Fe2O3 2Fe + Al2O3

24. Single Replacement Reactions • An active nonmetal can replace a less active nonmetal. • The halogen (F2, Cl2, Br2, I2) reactions are good examples. • F2 is the most active and I2 is the least. • Cl2 +2 NaI  2 NaCl + I2

25. Double Replacement Reactions • Ions of two compounds exchange places with each other. • The general form is AX + BY  AY + BX • Metathesis is an alternate name for double replacement reactions.

26. NaOH + CuSO4

27. Metathesis (sink or float?) • NaOH + CuSO4 • The Na+ and Cu2+ switch places. • Na+ combines with SO42- to form Na2SO4. • Cu2+ combines with OH- to form Cu(OH)2 • NaOH + CuSO4  Na2SO4 + Cu(OH)2 • 2NaOH + CuSO4  Na2SO4 + Cu(OH)2

28. CuSO4 + Na2CO3

29. Double Replacement • CuSO4 + Na2CO3 • Cu2+ combines with CO32- to form CuCO3. • Na+ combines with SO42- to form Na2SO4. • CuSO4 + Na2CO3  CuCO3 + Na2SO4

30. Na2CO3 + HCl

31. Double Replacement • Na2CO3 + HCl  • Notice that gas bubbles were produced rather than a precipitate. • What was the gas? • Write the double replacement reaction first.

32. Double Replacement • Na2CO3 + HCl  • Na+ combines with Cl- to form NaCl. • H+ combines with CO32- to form H2CO3. • Na2CO3 + 2HCl  2NaCl + H2CO3 • H2CO3 breaks up into H2O and CO2.

33. Double Replacement • The gas formed was carbon dioxide. • The final balanced reaction is: Na2CO3 + HCl  NaCl + H2O + CO2. • Balance the equation. • Na2CO3 + 2HCl  2NaCl + H2O + CO2

34. Combustion Reaction • When a substance combines with oxygen, a combustion reaction results. • The combustion reaction may also be an example of an earlier type such as 2Mg + O2 2MgO. • The combustion reaction may be burning of a fuel.

35. Combustion Reaction • Methane, CH4, is natural gas. • When hydrocarbon compounds are burned in oxygen, the products are water and carbon dioxide. • CH4 + O2 CO2 + H2O • CH4 + 2O2  CO2 + 2H2O

36. Combustion Reactions • Combustion reactions involve light and heat energy released. • Natural gas, propane, gasoline, etc. are burned to produce heat energy. • Most of these organic reactions produce water and carbon dioxide.

37. Practice • Classify each of the following as to type: • H2 + Cl2 2HCl • Combination • Ca + 2H2O  Ca(OH)2 + H2 • Single replacement

38. Practice • 2CO + O2 2CO2 • Combination and combustion • 2KClO3  2KCl + 3O2 • Decomposition

39. Practice • FeS + 2HCl  FeCl2 + H2S • Double replacement • Zn + HCl  ? • Single replacement • Zn + 2HCl  ZnCl2 + H2

40. Balancing Equations: Chemical

41. Ca O O H H How molecules are symbolized Cl2 2Cl 2Cl2 • Molecules may also have brackets to indicate numbers of atoms. E.g. Ca(OH)2 • Notice that the OH is a group • The 2 refers to both H and O • How many of each atom are in the following? • a) NaOH • b) Ca(OH)2 • c) 3Ca(OH)2 Na = 1, O = 1, H = 1 Ca = 1, O = 2, H = 2 Ca = 3, O = 6, H = 6

42. Mg Mg O O O  + Balancing equations: MgO • The law of conservation of mass states that matter can neither be created or destroyed • Thus, atoms are neither created or destroyed, only rearranged in a chemical reaction • Thus, the number of a particular atom is the same on both sides of a chemical equation • Example: Magnesium + Oxygen (from lab) • Mg + O2 MgO • However, this is not balanced • Left: Mg = 1, O = 2 • Right: Mg = 1, O = 1

43. Balance equations by “inspection” From Mg + O2 MgO 2Mg + O2 2MgO is correct Mg + ½O2 MgO is incorrect Mg2 + O2 2MgO is incorrect 4Mg + 2 O2 4MgO is incorrect Hints: start with elements that occur in one compound on each side. Treat polyatomic ions that repeat as if they were a single entity. 5 a) P4 + O2 P4O10 b) Li + H2O  H2+ LiOH c) Bi(NO3)3 + K2S Bi2S3 + KNO3 d) C2H6 + O2 CO2 + H2O 2 2 2 2 3 6 3.5 2 3 C2H6 + O2 CO2 + H2O 2 7 4 6

44. Balance these skeleton equations: a) Mg + 2HCl  MgCl2 + H2 b) 3Ca + N2 Ca3N2 c) NH4NO3 N2O + 2H2O d) 2BiCl3 + 3H2S  Bi2S3 + 6HCl e) 2C4H10 + 13O28CO2 + 10H2O f) 6O2 + C6H12O66CO2 + 6H2O g) 3NO2 + H2O 2HNO3 + NO h) Cr2(SO4)3+ 6NaOH  2Cr(OH)3+ 3Na2SO4 i) Al4C3 + 12H2O 3CH4 + 4Al(OH)3

45. The Mole Q: how long would it take to spend a mole of $1 coins if they were being spent at a rate of 1 billion per second?

46. Background: atomic masses • Look at the “atomic masses” on the periodic table. What do these represent? • E.g. the atomic mass of C is 12 (atomic # is 6) • We know there are 6 protons and 6 neutrons • Protons and neutrons have roughly the same mass. So, C weighs 12 u (atomic mass units). • What is the actual mass of a C atom? • Answer: approx. 2 x 10-23 grams (protons and neutrons each weigh about 1.7 x10-24 grams) Two problems • Atomic masses do not convert easily to grams • They can’t be weighed (they are too small)

47. The Mole Withtheseproblems,whyuseatomicmassatall? • Masses give information about # of p+, n0, e– • It is useful to know relative mass E.g. Q - What ratio is needed to make H2O? A - 2:1 by atoms, but 2:16 by mass • It is useful to associate atomic mass with a mass in grams. It has been found that 1gH,12gC,or 23gNahave6.02x1023atoms • 6.02 x 1023 is a “mole” or “Avogadro’s number” • “mol” is used in equations, “mole” is used in writing; one gram = 1 g, one mole = 1 mol.

48. Mollionaire Q: how long would it take to spend a mole of $1 coins if they were being spent at a rate of 1 billion per second? A: $ 6.02 x 1023 / $1 000 000 000 = 6.02 x 1014 payments = 6.02 x 1014 seconds 6.02 x 1014 seconds / 60 = 1.003 x 1013 minutes 1.003 x 1013 minutes / 60 = 1.672 x 1011 hours 1.672 x 1011 hours / 24 = 6.968 x 109 days 6.968 x 109 days / 365.25 = 1.908 x 107 years A: It would take 19 million years

49. Comparing sugar (C12H22O11) & H2O Same 1 gram each 1 mol each volume? No, they have dif. densities. No, molecules have dif. sizes. mass? Yes, that’s what grams are. No, molecules have dif. masses # of moles? No, they have dif. molar masses Yes. # of molecules? No, they have dif. molar masses Yes (6.02x1023 in each) # of atoms? No, sugar has more (45:3 ratio) No

50. Molar mass • The mass of one mole is called “molar mass” • E.g. 1 mol Li = 6.94 g Li • This is expressed as 6.94 g/mol • What are the following molar masses? S SO2 Cu3(BO3)2 32.06 g/mol 64.06 g/mol 308.27 g/mol Calculate molar masses (to 2 decimal places) CaCl2 (NH4)2CO3 O2 Pb3(PO4)2 C6H12O6 Cu x 3 = 63.55 x 3 = 190.65 B x 2 = 10.81 x 2 = 21.62 O x 6 = 16.00 x 6 = 96.00 308.27