1. Chapter 11 Chemical Reactions
Hingham High School
Mr. Dan Clune

2. All chemical reactions
Two parts:
Reactants – what you start with
Products- what you end with
Reactants turn into the products.
Reactants ® Products

3. The way atoms are joined is changed.
In a chemical reaction
The way atoms are joined is changed.
Atoms aren’t created or destroyed.

4. In a chemical reaction 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

5. Symbols in equations-p.323
Cu + Cl2  CuCl2
Arrow () separates reactants from products
Read “reacts to form”
Plus (+) sign read “and”

6. Symbols used in equations
Cu(s) + Cl2(g)  CuCl2(s)
(s) = solid
(g) = gas
(l) = liquid
(aq) - dissolved in water, an aqueous solution.

7. Symbols used in equations
­ ↑ after product, indicates gas produced
same as (g) - H2↑
¯ after product, indicates solid produced
same as (s) - PbI2↓

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

9. What is a catalyst? Speeds up a reaction
Is NOT changed or used up by the reaction.
Enzymes are biological or protein catalysts.

10. Skeleton Equation Uses formulas to describe a reaction
doesn’t indicate how many.
NOT balanced
All chem equations are description of rxn.

11. Convert this to an equation
Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas.
Fe2S3(s) + 
HCl(g)
FeCl3 +
H2S(g)

12. Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water, carbon dioxide gas, sodium nitrate dissolved in water. + 
HNO3(aq)
Na2CO3(aq) + +
H2O(l)
CO2(g)
NaNO3(aq)

13. 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)
Solid iron + oxygen gas yields solid iron (III) oxide.
Solid copper + aqueous silver nitrate yields solid silver + aqueous copper (II) nitrate.
Gaseous nitrite reacts with a platinum catalyst to form nitrogen gas and oxygen gas.

14. Balancing Chemical Equations

15. 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.

16. ® C + O2 ® CO2 This equation is already balanced What if it isn’t? O +
1 C and 2 O on each side.

17. ® C + O2 ® CO Need one more O in the products.
Can’t change the formula, because it describes what it is (carbon monoxide in this example)

18. ® Must be used to make another CO But where did the other C come from?+ C O C O
Must be used to make another CO
But where did the other C come from?

19. C C O O + O C C O
Must have started with two C
2 C + O2 ® 2 CO

20. Finding the number of atoms
The subscript in front of an element is the number of atoms of that element / polyatomic.
Ex) CO2 B2(SO4)3
C= O= B= SO4=
A coefficient in front of the formula multiplies the amount of elements by the coefficient.
Ex) 3CO2 2B2(SO4)3 1 2 2 3 3 6 4 6

21. Finding the number of atoms
H2O H= O=
2H2O
B(NO3)2 B=
NO3=
3B(NO3)2
Pb3(PO4)4
Pb=
PO4=
2Pb3(PO4)4 2 1 4 2 1 2 3 6 3 4 6 8

22. Rules for balancing:
. Determine the correct formulas for reactants and products.
. Write a skeleton equation.

23. Rules for balancing:
3. Count the number of atoms of each element appearing on both sides of the equation.
4. Balance the elements one at a time by adding coefficients (the numbers in front)
-Save H and O until LAST!

24. Rules for balancing: 5. Check to make sure it is balanced.
6. Make sure the coefficients are in the lowest possible ratio.

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

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

27. 2 2
H2 + O2
H2O

28. AgNO Cu ® Cu(NO3) Ag 2 2
NO3 Cu Ag P R 1 2 2 2 2

29. Mg + N2 ® Mg3N2 3 R P Mg N 1 2 3 3

30. P + O2 ® P4O10 4 5 R P O 1 2 4 10 4 10

31. Na + H2O ® H NaOH 2 2 2 O H Na P R 1 2 3 2 2 4 4 2 2

32. CH4 + O2 ® CO H2O 2 2 O H C P R 1 4 2 3 4 4 4

33. Section 11.2 Types of Chemical Reactions
OBJECTIVES:
Identify a reaction as combination, decomposition, single-replacement, double-replacement, or combustion

34. Types of Reactions 5 major types.
predict the products
predict whether or not they will happen at all
How? We recognize them by their reactants *

35. #1 - Combination Reactions
Combine - put together
2 or more substances combine to make one compound.
A + B ® AB
Ca +O2 ® CaO
SO3 + H2O ® H2SO4

36. #1 - Combination Reactions
We can predict the products if they are two elements.
Mg + N2 ® 3
Mg N 3 2

37. Write and balance Ca + Cl2 ® Fe + O2 ® iron (II) oxide Al + O2 ®
Remember that the first step is to write the correct formulas
Then balance by using coefficients only
CaCl2 2 2
FeO 4 3 2
Al2O3

38. #1 – Combination Reactions
Additional Notes:
a) Some nonmetal oxides react with H2O - produces acid:
SO2 + H2O  H2SO3
b) Some metallic oxides react with H2O - produces base:
CaO + H2O  Ca(OH)2
(how “acid rain” forms)

39. #2 - Decomposition Reactions
Decompose = fall apart
One reactant falls apart into two or more elements or compounds.
AB  A + B

40. #2 - Decomposition Reactions
NaCl Na + Cl2
CaCO CaO + CO2
Note that energy is usually required to decompose

41. #2 - Decomposition Reactions
Binary compounds (made of 2 elements) falls apart into its elements
H2O
HgO 2 2
H2 + O2 2
Hg + O2 2

42. #2 - Decomposition Reactions
If the compound has more than two elements, one product must be given.
The other product will be from the missing pieces
NiCO CO2 +
H2CO3(aq) ® CO2 + ? Ni H2

43. #3 - Single Replacement Reactions
One element replaces another (new dance partner)
Reactants are an element & cpd
Products will be a different element and different cmpd
Li + KCl ® K + LiCl
F2 +2 LiCl ® 2LiF + Cl2
(Cations switched)
(Anions switched) *

44. #3 Single Replacement Reactions
Metals replace other metals (and H)
K + AlN ®
Zn + HCl ®
Think of water as: HOH
Metals replace first H, then combines w/ hydroxide (OH).
Na H2O ® 3
K3N + Al 2
ZnCl H2 2 2
NaOH + H2 2 *

45. #3 Single Replacement Reactions
Sometimes, the reaction will not happen:
Some chemicals are more “active” than others
More active replaces less active *

46. The “Activity Series” of Metals
Lithium
Potassium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Nickel
Lead
Hydrogen
Bismuth
Copper
Mercury
Silver
Platinum
Gold
Higher activity
If the lone metal is above the paired metal, replacement will occur.
Ex) Li + NaCl ® Na + LiCl
If the lone metal is below the paired metal, replacement will not occur.
Ex) Na + LiCl ® Na + LiCl
The more active metal wants another element to stabilize it.
Lower activity

47. The “Activity Series” of Halogens
Higher Activity
If the lone halogen is above the paired halogen, replacement will occur.
Fluorine
Chlorine
Bromine
Iodine
Lower Activity
2NaCl(s) + F2(g) 
2NaF(s) + Cl2(g)
MgCl2(s) + Br2(g) 
No Reaction!

48. #3 Single Replacement Reactions Practice:
Lithium
Potassium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Nickel
Lead
Hydrogen
Bismuth
Copper
Mercury
Silver
Platinum
Gold
Higher activity
Fe + CuSO4 ®
Pb + KCl ®
Al + HCl ®
FeSO4 + Cu
No reaction 2 6 2
AlCl H2 3
Lower activity *

49. #4 - Double Replacement Reactions
Two things replace each other.
Reactants must be two ionic compounds.
NaOH + FeCl3 ®
positive ions change place
NaOH + FeCl3 ® Fe+3 OH- + Na+1 Cl-1 =
NaOH + FeCl3 ® Fe(OH)3 + NaCl *

50. #4 - Double Replacement Reactions
Have certain “driving forces”, or reasons
only happens if one product:
a) doesn’t dissolve in water & forms solid (a “precipitate”), or
b) is gas that bubbles out, or
c) is molecular compound (usually water) *

51. Complete and balance:
assume all of the following reactions actually take place:
CaCl NaOH ®
CuCl K2S ®
KOH + Fe(NO3)3 ®
K2SO BaF2 ® 2
Ca(OH) NaCl 2
CuS KCl 2 3 3
KNO Fe(OH)3 2
KF BaSO4 *

52. How to recognize which type?
Look at the reactants:
El + El = Combination
Cpd = Decomposition
El + Cpd = Single replacement
Cpd + Cpd = Double replacement *

53. Practice Examples: H2 + O2 ® H2O ® Zn + H2SO4 ® HgO ® KBr + Cl2 ®
AgNO NaCl ®
Mg(OH) H2SO3 ® 2
2H2O 2 2
H O2
H ZnSO4 2 2
Hg O2 2 2
KCl + Br2
AgCl + NaNO3
MgSO H2O 2 *

54. #5 – Combustion Reactions
Combustion means “add oxygen”
Normally, a cpd composed of only C, H, (sometimes O) is reacted with oxygen – called “burning”
Complete combustion, products are CO2 and H2O
If incomplete, products are CO and H2O *

55. Combustion Reaction Examples:
C4H O2 ®
C6H12O O2 ®
C8H O2 ® 7 4
CO H2O 6 9 6
CO H2O 6 10 8
CO H2O 4 *

56. SUMMARY: An equation... Describes a rxn Must be balanced
only balance by changing coefficients
special symbols to indicate physical state, catalyst or energy required, etc. *

57. Reactions
5 major types
We can tell what type they are by looking at reactants
Single Replacement happens based on the Activity Series *

58. Section 11.3 p. 342 Reactions in Aqueous Solution
KMnO4
Co(NO3)2
NiCl2
CuSO4
K2Cr2O7
K2CrO4

59. Net Ionic Equations Many reactions occur in water, or aqueous solution
When dissolved in water, many ionic cpds “dissociate”, or separate, into cations & anions
Now write ionic equation

60. Net Ionic Equations
Example (needs to be a double replacement reaction)
AgNO3 + NaCl  AgCl + NaNO3
1. this is the full balanced equation
2. next, write it as ionic equation by splitting the cpds into their ions:
Ag1+ + NO31- + Na1+ + Cl1- 
AgCl(s) + Na1+ + NO31-
Solids do not split up.

61. Net Ionic Equations
3. Crossing out ions that did not change (called spectator ions)
Ag1+ + Cl1-  AgCl (s)
This is the net ionic equation
Let’s talk about precipitates before we do some other examples

62. Predicting the Precipitate
Insoluble salt is a precipitate
i.e. a solid
General solubility rules are found:
Table 11.3, p. 344 in textbook

63. Solubility Rules

64. BaCl AgNO3 →
Break up into ions
Now write the net ionic equation. 2
Ba(NO3) AgCl (s) 2
Ba+2 + 2Cl- + 2Ag+ + 2NO3-  Ba+2 + 2NO3- + 2AgCl (s)
2Ag+ + 2NO3-  2AgCl (s)

65. Now write the net ionic equation. 2
NaCl + Ba(NO3)2 → 2
NaNO BaCl2
Break up into ions
Now write the net ionic equation.
2Na+ + 2Cl- + Ba+2 + 2NO3-  2Na+ + 2NO3- + Ba+2 + 2Cl-
Nothing changed, so no net ionic equation.

66. PbCl2 (s) + Li2O 
PbO LiCl 2
Break up into ions
Now write the net ionic equation.
PbCl2 (s) + 2Li+ + O-2  Pb+2 + O-2 + 2Li+ + 2Cl-
PbCl2 (s)  Pb+2 + 2Cl-