1. Unit 5 Chemical Equations & Reactions
Honors Chemistry

2. All chemical reactions
have two parts:
Reactants - the substances you start with
Products- the substances you end up with
The reactants turn into the products.
Reactants ® Products

3. In a chemical reaction The way atoms are joined is changed
Atoms aren’t created of destroyed.
Can be described several ways:
1. In a sentence
Copper reacts with chlorine to form copper (II) chloride.
2. In a word equation
Copper + chlorine ® copper (II) chloride

4. Symbols in equations
the arrow separates the reactants from the products
Read “reacts to form”
The plus sign = “and”
(s) after the formula = solid
(g) after the formula = gas
(l) after the formula = liquid

5. Symbols used in equations
(aq) after the formula - dissolved in water, an aqueous solution.
­ used after a product indicates a gas (same as (g))
¯ used after a product indicates a solid (same as (s))

6. Symbols used in equations
indicates a reversible reaction
shows that heat is supplied to the reaction
is used to indicate a catalyst is supplied, in this case, platinum.

7. What is a catalyst?
A substance that speeds up a reaction, without being changed or used up by the reaction.
Enzymes are biological or protein catalysts.

8. Skeleton Equation Uses formulas and symbols to describe a reaction
doesn’t indicate how many.
All chemical equations are sentences that describe reactions.

9. Convert these to equations
Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas.
Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

10. Now, read these: Fe(s) + O2(g) ® Fe2O3(s)
Cu(s) + AgNO3(aq) ® Ag(s) + Cu(NO3)2(aq)
NO2 (g) N2(g) + O2(g)

11. Balanced Equation Atoms can’t be created or destroyed
All the atoms we start with we must end up with
A balanced equation has the same number of each element on both sides of the equation.

12. Rules for balancing:
Assemble, write the correct formulas for all the reactants and products
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) - save H and O until LAST!
Check to make sure it is balanced.

13. Never Never change a subscript to balance an equation.
If you change the formula you are describing a different reaction.
H2O is a different compound than H2O2
Never put a coefficient in the middle of a formula
2 NaCl is okay, Na2Cl is not.

14. Example
H2 + O2
H2O R P 2 H 2 2 O 1
Need twice as much O in the product

15. Example
H2 + O2 2
H2O R P 2 H 2 2 O 1
Changes the O

16. Example
H2 + O2 2
H2O R P 2 H 2 2 O 1 2
Also changes the H

17. Example
H2 + O2 2
H2O R P 2 H 2 4 2 O 1 2
Need twice as much H in the reactant

18. Example 2
H2 + O2 2
H2O R P 2 H 2 4 2 O 1 2
Recount

19. Example 2
H2 + O2 2
H2O R P 4 2 H 2 4 2 O 1 2
The equation is balanced, has the same
number of each kind of atom on both sides

20. Example R P 4 2 H 2 4 2 O 1 2 2 H2 + O2 ® 2 H2O This is the answer
Not this

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

22. Balancing Examples 2AgNO3 + Cu ® Cu(NO3)2 + 2Ag 3Mg +N2 ® Mg3N2
4P + 5O2 ® P4O10
2Na + 2H2O ® H2 + 2NaOH
CH4 + 2O2 ® CO2 + 2H2O

23. Types of Reactions There are millions of reactions.
Can’t remember them all
Fall into several categories.
We will learn 5 major types.
Will be able to predict the products.
For some, we will be able to predict whether they will happen at all.
Will recognize them by the reactants

24. CaO(s) + SO2(g)  CaSO3(s)
COMBINATION REACTION
A reaction in which two or more substances combine to form a single product.
A +B + C  ABC
CaO(s) + SO2(g)  CaSO3(s)

25. Write and balance Ca + Cl2 ® Fe + O2 ® Al + O2 ®
Remember that the first step is to write the correct formulas
Then balance by using coefficients only

26. DECOMPOSITION REACTION
A reaction in which a single compound reacts to give two or more substances, usually requiring a raise in temperature.
ABC  A + B + C
2KClO3(s)  2KCl(s) + 3O2(g)

27. #2 - Decomposition Reactions
decompose = fall apart
one reactant falls apart into two or more elements or compounds.
NaCl Na + Cl2
CaCO CaO + CO2
Note that energy is usually required to decompose

28. #2 - Decomposition Reactions
Can predict the products if it is a binary compound
Made up of only two elements
Falls apart into its elements
H2O
HgO

29. Can predict the products if it is a binary compound
2H2O(l) →
Can predict the products if it is a binary compound
Made up of only two elements
Falls apart into its elements
Fig. 8-13, p. 215

30. #2 - Decomposition Reactions
If the compound has more than two elements you must be familiar with page 10 of your packet
NiCO3 ®
KClO3(aq) ®

31. Single Displacement (or Replacement) Reactions
pp. 218, 220

32. Fig. 8-15, p. 218

33. Single Displacement (or Replacement) Reactions
PREDICT THE PRODUCT
1. Ca + HCl 
2. ZnBr2 + I2 
3. Cu + AgNO3 
Answers are on the next slide.

34. #3 Single Replacement Metals replace other metals (and hydrogen)
K + AlN ®
Zn + HCl ®
Think of water as HOH
Metals replace one of the H, combine with hydroxide.
Na + HOH ®

35. We can tell whether a reaction will happen
Some chemicals are more “active” than others
More active replaces less active
There is a list on page 11 - called the Activity Series of Metals
Higher on the list replaces lower.

36. #3 Single Replacement Note the * concerning Hydrogen
H can be replaced in acids by everything higher
Li, K, Ba, Ca, & Na replace H from acids and water
Fe + CuSO4 ®
Pb + KCl ®
Al + HCl ®

37. #3 - Single Replacement
What does it mean that Hg and Ag are on the bottom of the list?
Nonmetals can replace other nonmetals
Limited to F2 , Cl2 , Br2 , I2 (halogens)
Higher replaces lower.
F2 + HCl ®
Br2 + KCl ®

38. Double Displacement (or Replacement) Reactions
pp. 220, 223

39. Double Displacement (or Replacement) Reactions
PREDICT THE PRODUCT & BALANCE
1. MgSO4 + LiOH  ___________
2. Pb(NO3)2 + Na2CO3  ____________
3. HNO3 + Ba(OH)2  ___________
Answers are on the next slide.

40. Double Displacement (or Replacement) Reactions
ANSWERS
1. MgSO4 + 2 LiOH  Mg(OH)2 + Li2SO4
2. Pb(NO3)2 + Na2CO3  PbCO3 + 2 NaNO3
3. 2 HNO3 + Ba(OH)2  Ba(NO3)2 + 2 H2O
Exchange cations

41. Double Replacement Two things replace each other.
Reactants must be two ionic compounds or acids.
Usually in aqueous solution
NaOH + FeCl3 ®
The positive ions change place.
NaOH + FeCl3 ® Fe+3 OH- + Na+1 Cl-1
NaOH + FeCl3 ® Fe(OH)3 + NaCl

42. Double Replacement Has certain “driving forces”
Will only happen if one of the products:
doesn’t dissolve in water and forms a solid (a “precipitate”), or
is a gas that bubbles out, or
is a covalent compound (usually water).

43. Oxidation and Reduction Reactions
Early chemists saw oxidation as the combination of a material with oxygen
Today, many of these reactions may not even involve oxygen
Redox currently says that electrons are transferred between reactants
Always occur simultaneously.
Called redox reactions.

44. Oxidation and Reduction Reactions
Oxidation – refers to a loss of electrons, or gain of oxygen
Reduction – refers to a gain of electrons, or loss of oxygen
Oxidizing agent – the substance being reduced
Reducing agent – the substance being oxidized

45. Oxidation and Reduction Reactions
Loss of Electrons is Oxidation
Gain of Electrons is Reduction
LEO SAYS GER

46. Rules For Assigning Oxidation States
Oxidation numbers are used to keep track of how many electrons are lost or gained by each atom

47. Balancing Redox Reactions
If the oxidation number changes, that element has undergone oxidation or reduction
Reaction as a whole must be redox
Half-reaction – equation showing just the oxidation or just the reduction

48. Balancing Redox Reactions (in an acid)
SO2 + MnO4- →SO42- + Mn2+
Divide the equation into half-reactions.
SO2 → SO MnO4- → Mn2+
Oxidation Reduction
Balance all atoms except O and H

49. Balancing Redox Equations
Balance O by adding H2O.
2H2O + SO2 →SO MnO4- →Mn2+ + 4H2O
Balance H by adding H+.
2H2O + SO2 → SO H+
8H+ + MnO4- → Mn2+ + 4H2O

50. Balance net charge by adding e-.
2H2O + SO2 →SO H+ + 2e-
5e- + 8H+ + MnO4- →Mn2+ + 4H2O Step 6.
Make e- gain equal e- loss; then add half- reactions.
5(2H2O + SO2 →SO H+ + 2e-) 2(5e- + 8H+ + MnO4- →Mn2+ + 4H2O) ______________________________________ 10H2O + 5SO2 + 10e- + 16H+ + 2MnO4- →2Mn2+ + 8H2O + 5SO H+ + 10e-

51. Cancel anything that's the same on both sides.
2H2O + 5SO2 +2MnO4- →2Mn2+ + 5SO H+

52. Balancing Redox Reactions (in a base)
Follow steps 1 – 7 then: 8. Identify the number of proton (H+) in the acidic answer. Add the same number of OH+ ions to BOTH sides of the equation. 9. If H+ and OH+ appear on the same side of the equation, they will react in a 1:1 ratio to form H2O. 10. Cancel out water molecules that appear on both sides of the equation. 11. Check that all is balanced

53. PRECIPITATION REACTION
A reaction where an insoluble solid is formed during a reaction between two aqueous solutions.
(aq) (aq)  (aq) (s)
2KI(aq) + Pb(NO3)2(aq)  2KNO3(aq) + PbI2(s)
Exchange cations

54. Ionic Equations
Many reactions occur in water- that is, in aqueous solution
Many ionic compounds “dissociate”, or separate, into cations and anions when dissolved in water
AgNO3 + NaCl  AgCl + NaNO3
1. this is the full equation
2. Write it as an ionic equation
3. Eliminating ions not directly involved (spectator ions) = net ionic equation

55. Molecular, Ionic, Net Ionic Equations

56. How Do We Predict Solubility?

57. NEUTRALIZATION REACTION
A reaction between an acid and a base which results in the production of a salt and water.
HA + BOH  (metal/nonmetal) + H2O
HNO3(aq) + KOH(aq)  KNO3(aq) + H2O(l)

58. 2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g)
COMBUSTION REACTION
Means “add oxygen”
A compound composed of only C, H, and maybe O is reacted with oxygen
If the combustion is complete, the products will be CO2 and H2O.
If the combustion is incomplete, the products will be CO (possibly just C) and H2O.
A reaction of a substance with oxygen, usually the rapid release of heat produces a flame.
CH + O2  CO2 + H2O
2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g)

59. 2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g) + heat (in Joules)
Combustion
Many times in a combustion reaction, heat energy is given off. In chemical terms this is called an exothermic reaction.
Thermochemistry is field of chemistry which studies the transfer of heat in a reaction.
The thermodynamic equation representing this exothermic reaction is:
2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g) + heat (in Joules)

60. Examples C4H10 + O2 ® (assume complete) C4H10 + O2 ® (incomplete)
C6H12O6 + O2 ® (complete)
C8H8 +O2 ® (incomplete)

61. GAS FORMATION REACTIONS
A reaction that produces a gas from reactants not in the gaseous state.
2 HCl (aq) + ZnS (s)  ZnCl2 (aq) + H2S (g)
Zn (s) + 2 HCl (aq)  ZnCl2 (aq) + H2 (g)
Many gas formation reactions involve two steps, first the double displacement reaction then the decomposition reaction of an unstable substance.
Na2CO3 + 2HCl  2 NaCl + H2CO3
H2CO3  CO2 + H2O
Besides carbonic acid (H2CO3), sulfurous acid (H2SO3) also decomposes into SO2 and water.

62. COMMON GAS FORMATION REACTIONS YOU SHOULD REMEMBER
NH4OH → NH3 (g) + H2O (l)
H2CO3 → CO2 (g) + H2O (l)
H2SO3 → SO2 (g) + H2O (l)

63. Remember an equation... Describes a reaction
Must be balanced in order to follow the Law of Conservation of Mass
Can only be balanced by changing the coefficients.
Has special symbols to indicate physical state, and if a catalyst or energy is required.

64. Table 8-3, p. 223

65. Examples H2 + O2 ® H2O ® Zn + H2SO4 ® HgO ® KBr +Cl2 ® AgNO3 + NaCl ®
Mg(OH)2 + H2SO3 ®