1. Chapter 8 Chemical Reactions

2. Section 1: Objectives
Identify the parts of a chemical equation
Learn how to write a chemical equation
Learn how to balance a chemical equation.

3. A chemical reaction is the process
by which one or more substances are
changed into one or more different
substances.
Chemical reactions are described by chemical equations.
A chemical equations represents, with symbols and formulas, the identities and relative amounts of the starting material and products.

4. Indications of a Chemical Reaction
Evolution of heat and light.
Production of a gas.
Formation of a precipitate (solid).
(a solid appears after two liquids are mixed)
4) Color change

5. All chemical reactions…
have two parts:
Reactants = the substances you start with (left side of arrow)
Products = the substances you end up with (right side of arrow)
The reactants will turn into the products.
Reactants ® Products

6. In a chemical reaction
Atoms aren’t created or destroyed (according to the Law of Conservation of Mass)
A reaction can be described several ways:
#1. In a sentence every item is a word:
Copper reacts with chlorine to form
copper (II) chloride.
#2. In a word equation some symbols used:
Copper + chlorine ® copper (II) chloride

7. Symbols in equations
the arrow (→) separates the reactants from the products (arrow points to products)
Read as: “reacts to form” or yields
(s) after the formula = solid: Fe(s)
(g) after the formula = gas: CO2(g)
(l) after the formula = liquid: H2O(l)
(aq) after the formula = dissolved in water, an aqueous solution: NaCl(aq) is a salt water solution

8. Symbols in equations
­ ↑ used after a product indicates a gas has been produced: H2↑
¯ used after a product indicates a solid has been produced: PbI2↓

9. Symbols in equations shows that heat is supplied to the reaction
Additional symbols on page 246 of textbook.

10. Formula Equations:
Represents the reactants and products by symbols and formulas:
Fe(s) + O2(g) ® Fe2O3(s)
Cu(s) + AgNO3(aq) ® Ag(s) + Cu(NO3)2(aq)
NO2(g) ® N2(g) + O2(g)

11. Balanced Chemical Equations
Atoms can’t be created or destroyed in an ordinary reaction:
All the atoms we start with we must end up with (meaning: balanced!)
A balanced equation has the same number of each element on both sides of the equation.

12. Rules for balancing:
Assemble the correct formulas for all the reactants and products, using “+” and “→”
Count the number of atoms of each type appearing on both sides
Balance the elements one at a time by adding coefficients (the numbers in front)
(order of balancing on next slide)
Double-Check to make sure it is balanced.

13. Rules for balancing: Metals Non-metals
save balancing the H and O until LAST!
(hint: I prefer to save O until the very last)

14. Rules
Never change a subscript to balance an equation (You can only change coefficients)
If you change the subscript (formula) you are describing a different chemical.
H2O is a different compound than H2O2
Never put a coefficient in the middle of a formula; they must go only in the front
2NaCl is okay, but Na2Cl is not.

15. Nitrogen + hydrogen ammonia
Example 1
N2 + H NH3
Nitrogen + hydrogen ammonia
Count atoms.
Reactants: 2 atoms N and 2 atoms H
Products: 1 atom N and 3 atoms of NH3

16. Balancing Equations N2 + H2 2NH3 Nothing is balanced.
Balance the nitrogen first by placing a coefficient of 2 in front of the NH3.
N2 + H NH3

17. Balancing Equations N2 + 3H2 2NH3 Hydrogen is not balanced.
Place a 3 in front of H2.
Reactant side: 2 atoms N, 6 atoms H
Product side: 2 atoms N, 6 atoms H
It’s balanced!
N H NH3

18. Example 2 Cu + H2SO4 CuSO4 + H2O + SO2 Count Atoms:
Reactants: Cu – 1, H – 2, S – 1, O – 4
Products: Cu – 1, H – 2, S - 2, O - 7

19. Balancing Equations Cu + 2H2SO4 CuSO4 + H2O + SO2
Sulfur is not balanced.
Place a two in front of sulfuric acid.
Count atoms: 2 H2SO4 H – 4, S – 2, O - 8
Cu + 2H2SO4 CuSO4 + H2O + SO2

20. Balancing Equations Cu + 2H2SO4 CuSO4 + 2H2O + SO2
Hydrogen needs to be balanced so place a 2 in front of the H2O.
Count the number of atoms.
Cu + 2H2SO4
CuSO H2O + SO2

21. Balancing Equations Cu + 2H2SO4 CuSO4 + 2H2O + SO2
Reactants: Cu – 1, H – 4, S – 2, O – 8
Products: Cu – 1, H – 4, S – 2, O – 8
It’s balanced!
Cu + 2H2SO4
CuSO H2O + SO2

22. Practice Balancing Examples
_AgNO3 + _Cu ® _Cu(NO3)2 + _Ag
_Mg + _N2 ® _Mg3N2
_P + _O2 ® _P4O10
_Na + _H2O ® _H2 + _NaOH
_CH4 + _O2 ® _CO2 + _H2O 2 2 3 4 5 2 2 2 2 2

23. Section 2 Types of Chemical Reactions
OBJECTIVES:
Describe the five general types of reactions.
Predict the products of the five general types of reactions.

24. Type of Reactions
Chemical reactions are classified into five general types
Combination (synthesis)
Decomposition
Single Replacement
Double Replacement
Combustion

25. Combination (Synthesis)
Two or more elements or simple compounds combine to form (synthesize) one product
A + B AB

26. Combination (Synthesis)
1) Reactions of elements with oxygen and sulfur.
The groups 1 and 2 metals react with oxygen to form oxides.
2 Mg + O MgO (group 2)
4 Li + O2 2 Li2O (group 1)

27. Combination (Synthesis)
1) Reactions with oxygen and sulfur.
Sulfur which is right below oxygen reacts in a similar manner
8 Ba + S BaS
16 Rb + S8 8 Rb2S

28. Combination (Synthesis)
2) Reactions of metals with halogens
2 Na + Cl NaCl
2 K + I KI
Mg + F MgF2

29. Decomposition
One substance is broken down (split) into two or more simpler substances.
These reactions usually require heat
AB A + B

30. Decomposition 1) Decomposition of binary compounds
2 H2O 2 H2 (g) + O2 (g) (electrolysis)
electricity
2 HgO 2 Hg + O2

31. Decomposition 2) Decomposition of metal carbonates CaCO3 CaO + CO2
3) Decomposition of metal hydroxides
Ca(OH)2 CaO + H2O

32. Decomposition 4) Decomposition of metal chlorates 2 KClO3 2 KCl + 3 O2
5) Decomposition of acids
H2CO3 CO2 + H2O (lab)

33. Learning Check Classify the following reactions as
1) combination or 2) decomposition:
___A. H2 + Br HBr
___B. Al2(CO3)3 Al2O3 + 3CO2
___C. 4 Al + 3C Al4C3

34. Solution Classify the following reactions as
1) combination or 2) decomposition:
_1_A. H2 + Br HBr
_2_B. Al2(CO3)3 Al2O3 + 3CO2
_1_C. 4 Al + 3C Al4C3

35. Single Replacement
One element replaces a similar element in a compound.
A + BX AX + B
Zn + 2HCl ZnCl H2
Fe CuSO4 FeSO4 + Cu
Many single replacement reactions take place in an aqueous solution.

36. Reaction is not balanced.
Single Replacement
1) Replacement of a metal by another metal.
Refer to activity series table to determine if the reaction will occur.
Al + Pb(NO3) Pb + Al(NO3)3
Reaction is not balanced.
Try to balance.

37. Single Replacement 1) Replacement of a metal by another metal.
Refer to activity series table to determine if the reaction will occur.
2 Al Pb(NO3) Pb + 2 Al(NO3)3
Balanced!

38. Single Replacement 2) Replacement of hydrogen in water by a metal.
Na + H2O NaOH + H2 (g)
Unbalanced!
Group 1 metals react vigorously with water.

39. Single Replacement 2) Replacement of hydrogen in water by a metal.
2 Na H2O NaOH + H2 (g)
Balanced!

40. Single Replacement 3) Replacement of hydrogen in an acid by a metal.
Mg HCl MgCl2 + H2 (g)
balanced!
Group 1 and 2 metals can react with acids to produce hydrogen gas.

41. Single Replacement 4) Replacements of halogens. Cl2 + KBr KCl + Br2
unbalanced!
Refer to activity series table.
Order: F > Cl > Br > I

42. Single Replacement 4) Replacements of halogens.
Cl KBr KCl + Br2
balanced!
Refer to activity series table.

43. Single Replacement 4) Replacements of halogens.
F NaCl NaF + Cl2
balanced!
Refer to activity series table.
Order: F > Cl > Br > I

44. Single Replacement 4) Replacements of halogens. Br2 + KCl products?
Refer to activity series table.
Order: F > Cl > Br > I

45. Double Replacement Two elements in reactants exchange places.
AX + BY AY + BX
AgNO3 + NaCl AgCl + NaNO3
ZnS HCl ZnCl2 + H2S
One new product is usually a precipitate or gas

46. Double Replacement 1) Formation of a precipitate.
KI + Pb(NO3)2 KNO PbI2 (s)
unbalanced
Exchange of metals in the reaction.
PbI2 is a solid that precipitates out of
the reaction.

47. Double Replacement 1) Formation of a precipitate.
2 KI + Pb(NO3)2 2 KNO PbI2 (s)
balanced
Exchange of metals in the reaction.

48. Iron(II) sulfide reacts to form
Double Replacement
2) Formation of a gas.
Fe(II)S HCl H2S (gas) + FeCl2
balanced
Iron(II) sulfide reacts to form
hydrogen sulfide gas.

49. Usually an acid-base reaction
Double Replacement
3) Formation of a water.
HCl + NaOH NaCl + H2O (liquid)
balanced
Usually an acid-base reaction

50. Learning Check Classify the following reactions as
1) single replacement
2) double replacement
__A) 2Al + 3H2SO Al2(SO4) H2
__B) Na2SO4 + 2AgNO Ag2SO4 + 2NaNO3

51. Solution Classify the following reactions as 1) single replacement
2) double replacement
1_A) 2Al + 3H2SO Al2(SO4) H2
2_B) Na2SO4 + 2AgNO Ag2SO4 + 2NaNO3

52. Combustion
A reaction in which a compound (often carbon) reacts with oxygen.
C + O2 CO2
CH4 + 2 O2 CO H2O
C3H O2 3 CO2 + 4 H2O
C6H12O O CO2 + 6 H2O
Reaction releases energy in the form of heat or light.

53. Learning Check ___C5H12 + ___O2 ___CO2 + ___H2O
Balance the combustion equation
___C5H ___O ___CO2 + ___H2O

54. Solution
Balance the combustion equation
1 C5H O CO H2O

55. Learning Check ___C6H14 + ___O2 ___CO2 + ___H2O
Balance the combustion equation
___C6H ___O ___CO2 + ___H2O

56. Solution 1 C6H14 + 9.5 O2 6 CO2 + 7 H2O Fractions are not allowed!
Balance the combustion equation
1 C6H O CO H2O
Fractions are not allowed!
Solution: double all the coefficients

57. Solution
Balance the combustion equation
2 C6H O CO H2O
Solution: double all the coefficients

58. End

59. Oxidation and Reduction
Reactions that involve a loss or gain of electrons
Occurs in many of the 4 types of reactions and combustion
Important in food metabolism, batteries, rusting of metals

60. Requirements for Oxidization-Reduction
Electrons are transferred
Two processes occur
Oxidation = Loss of electrons (LEO)
Zn Zn2+ + 2e-
Reduction = Gain of electrons (GER)
Cu e Cu

61. Balanced Red-Ox Equations
Combine the oxidation and reduction reactions to make
Loss of electrons = Gain of electrons
Zn + Cu2+ + 2e- Zn2+ + 2e- + Cu
Zn + Cu2+ Zn2+ + Cu

62. Gain/Loss of Hydrogen In organic and biological reactions
oxidation = Loss of H / Gain of Oxygen
reduction = Gain of H / Loss of oxygen

63. Learning Check R3
Identify the following as an 1) oxidation or a reduction process:
__A. Sn Sn e-
__B. Fe e- Fe2+
__C. Cl2 + 2e- 2Cl-

64. Solution R3
Identify the following as an 1) oxidation or a reduction process:
1_ A. Sn Sn e-
2_ B. Fe e- Fe2+
2_ C. Cl2 + 2e- 2Cl-

65. Learning Check R4
In light-sensitive sunglasses, UV light initiates an oxidation-reduction reaction
Ag+ + Cl Ag Cl
A. Which reactant is oxidized
1) Ag ) Cl ) Ag
B. Which reactant is reduced?
1) Ag ) Cl ) Cl

66. Solution R4
In light-sensitive sunglasses, UV light initiates an oxidation-reduction reaction
Ag+ + Cl Ag Cl
A. Which reactant is oxidized
2) Cl Cl- Cl + e-
B. Which reactant is reduced?
1) Ag+ Ag+ + e Ag