1. Chemistry of Matches P4S3 + KClO3 P2O5 + KCl + SO2 D tetraphosphorus
trisulfide
potassium
chlorate
diphosphorus
pentoxide
potassium
chloride
sulfur
dioxide
                                                   
Strike anywhere matches
Safety matches
The substances P4S3 and KClO3 are present on the tip of a match. When the
match is struck on a rough surface, the two chemicals (reactants) ignite and
produce a flame. The products from this reaction are P2O5, KCl, and SO2,
the last which is responsible for the characteristic sulfur smell.
Safety matches have the P4S3 on the matchbook or box. This is the red strip you need to rub the match against to get it to ignite.
Strike anywhere matches contain both reactants on the match head.
[NOTE: Red phosphorus is an important ingredient for making methamphetamines – and may be reason high school teachers are not allowed to purchase it.]
A good demonstration is the SMOKE reaction: combine 6 g KClO3 + 2 g table sugar (stir gently) and place in an evaporating dish.
Add 1 drop of [H2SO4]. The sugar will rapidly burn! Do this reaction only in a fume hood. It will produce lots of smoke. The rapid
combustion of the sugar is due to the fact that the potassium chlorate is an oxidizer (contains oxygen).
Oxygen is not the limiting reactant for this demonstration.
I also show the DEHYDRATION OF SUGAR demonstration. I ½ fill a 250 mL beaker with table sugar, add ~5 mL tap water and stir with a glass stirring rod. Finally, add ~50 mL concentrated sulfuric acid. The sugar will be rapidly dehydrated (leaving only a carbon tower). DO NOT touch the burned sugar as it contains excess sulfuric acid.
This reaction is a much slower burning of table sugar…not enough oxygen present.
The substances P4S3 and KClO3 are both
present on the tip of a strike anywhere match.
When the match is struck on a rough surface,
the two chemicals (reactants) ignite and produce a flame.
The substances P4S3 and KClO3
are separated. The P4S3 is on
the matchbox cover.
Only when the chemicals combine
do they react and produce a flame.
The products from this reaction are P2O5, KCl, and SO2,the
last of which is responsible for the characteristic sulfur smell.
Charles H.Corwin, Introductory Chemistry 2005, page 182

2. Chemical Equations and Reactions
Chapter 8

3. Describing Chemical Reactions
Section 1

4. (NH4)2Cr2O7 (s)  N2(g) + Cr2O3(s) + 4H2O(g)
Chemical reaction  process by which one or more substances are changed into one or more different substances
Represented by chemical equations  represents, with symbols and formulas, the identities and relative amounts of the reactants and products in a chemical reaction
(NH4)2Cr2O7 (s)  N2(g) + Cr2O3(s) + 4H2O(g)

5. Organize Your Thoughts
Chemical
reactions
Chemical
equations
Chemical
equations
To identify the reactants and products in a chemical equation.
To balance a chemical equation with coefficients.
To identify synthesis (combination), decomposition, and combustion reactions.
To identify single-replacement reactions.
To identify double-replacement reactions.
To predict the products of a reaction.
Synthesis
Decomposition
Single replacement
Double replacement
Combustion
Balancing equations
Predicting products
from reactants
Packard, Jacobs, Marshall, Chemistry Pearson AGS Globe, page 175

6. Signs of Chemical Reactions
There are five main signs that indicate a chemical reaction has taken place:
release
input
change in color
change in odor
production of new
gases or vapor
input or release
of energy
difficult to reverse

7. Chemical Equations Depict the kind of reactants and products
aluminum oxide
Depict the kind of reactants and products
and their relative amounts in a reaction.
reactants
product
4 Al(s) O2(g) Al2O3(s)
A chemical equation is an expression that gives the identities and quantities of the substances in a chemical reaction
Chemical formulas and other symbols are used to indicate the starting material(s) or reactant(s), which are written on the left side of the equation, and the final compound(s) or product(s), which are written on the right side. An arrow, read as yields or reacts to form, points from the reactants to the products.
Abbreviations are added in parentheses as subscripts to indicate the physical state of each species:—(s) for solid, (l) for liquid, (g) for gas, and (aq) for an aqueous solution.
A balanced chemical equation is when both the numbers of each type of atom and the total charge are the same on both sides. A chemical reaction represents a change in the distribution of atoms but not in the number of atoms.
The letters (s), (g), and (l) are the
physical states of compounds.
The numbers in the front are called
stoichiometric coefficients.

8. Chemical Equations 4 Al(s) + 3 O2(g) 2 Al2O3(s)
aluminum oxide
sandpaper
4 Al(s) O2(g) Al2O3(s)
4 g Al g O2 yield 2 g Al2O3
This equation means:
4 Al atoms O2 molecules yield 2 molecules of Al2O3 or
4 Al moles O2 moles yield 2 moles of Al2O3
4 mol
3 mol
2 mol
108 g g = g

9. CH4(g) + O2(g)  CO2(g) + H2O(g)
Formula equations
Formula equation  represents reactants and products of chemical reaction by their symbols or formulas
CH4(g) + O2(g)  CO2(g) + H2O(g)
Is the law of conservation of mass satisfied here?

10. CH4 + O2  CO2 + H2O Reactants – 7 atoms Products – 6 atoms
Hydrogen = 4 on left, = 2 on right
Oxygen = 2 on left, = 3 on right
In order to satisfy law of conservation of mass, must balance the equation by adding coefficients

11. Balancing a Chemical Equation
Write a formula equation by substituting correct formulas for the names of the reactants and products (if you do not start with a formula equation).
H2O(l)  H2(g) + O2(g)
Balance the formula equation according to the law of conservation of mass. There are a few simple guidelines to use for this…..

12. Guidelines for Balancing Equations
Balance the different types of atoms one at a time.
First balance the atoms of elements that are combined and that appear only once on each side of the equation.
Balance polyatomic ions that appear on both sides as single units.
Balance H and O atoms (or any other lone atoms) last.

13. 2H2O(l)  2H2(g) + O2(g) ---- balanced!
2 oxygen on right and only 1 on the left
Start with placing 2 in front of H2O
2H2O(l)  H2(g) + O2(g)
Oxygen is now balance with 2 on left and 2 on right
Now balance hydrogen – 4 on left, 2 on right
Add coefficient 2 to H2
2H2O(l)  2H2(g) + O2(g) balanced!

14. DO NOT!.....
DO NOT WRITE INCORRECT FORMULAS, THIS WILL MESS UP YOUR BALANCING
DO NOT CHANGE SUBSCRIPTS IN FORMULAS TO BALANCE THE EQUATION!
H2O(l)  H2(g) + O2(g) H2O(l)  H2(g) + O(g)

15. DO!!!!.....
When you think you have balanced the equation, COUNT THE NUMBERS OF EACH TYPE OF ATOM ON EITHER SIDE OF THE EQUATION
2H2O(l)  2H2(g) + O2(g)
Reactant side
Product side H 4 O 2

16. Sample Problem
The reaction of zinc with aqueous hydrochloric acid produces a solution of zinc chloride and hydrogen gas. Write a balanced equation for the reaction.

17. 1. Write the word equation
Zinc + hydrochloric acid  zinc chloride + hydrogen
2. Write the formula equation
Zn(s) + HCl(aq)  ZnCl2(aq) + H2(g)

18. Adjust the coefficients
Zn(s) + HCl(aq)  ZnCl2(aq) + H2(g) Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) Count atoms: Zn – 1 on left, 1 on right H – 2 on left, 2 on right Cl – 2 on left, 2 on right

19. Practice Problem 1a
Write word, formula, and balanced chemical equations for magnesium and hydrochloric acid react to produce magnesium chloride and hydrogen.
Magnesium + hydrochloric acid  magnesium chloride + hydrogen
Mg + HCl  MgCl2 + H2
Mg + 2HCl  MgCl2 + H2

20. Practice Problem 1b
Write word, formula, and balanced chemical equations for silicon dioxide and hydrofluoric acid reacting to produce silicon tetrafluoride and water.
Silicon dioxide+ hydrofluoric acid  silicon tetrafluoride + water
SiO2+ HF  SiF4 + H2O
SiO2+ 4HF  SiF4 + 2H2O

21. Practice Problem 2
Write word, formula and balanced equations for aqueous nitric acid reacts with solid magnesium hydroxide to produce aqueous magnesium nitrate and water.
Nitric acid + magnesium hydroxide  magnesium nitrate + water
HNO3(aq) + Mg(OH)2(s)  Mg(NO3)2(aq) + H2O(l)
2HNO3(aq) + Mg(OH)2(s)  Mg(NO3)2(aq) + 2H2O(l)

22. Practice Problem 3
Ammonium sulfate crystals are made by treating ammonia gas, often a by-product from coke-ovens, with aqueous sulfuric acid:
2 NH3(g) + H2SO4(aq) → (NH4)2SO4(s)

24. Practice Problem 4
Aluminum sulfate and calcium hydroxide are used in a water-purification process. When added to water, they dissolve and react to produce 2 insoluble products, aluminum hydroxide and calcium sulfate. Write a balanced equation for this reaction.
Al2(SO4)3(aq) + 3Ca(OH)2(aq)  2Al(OH)3(s) + 3CaSO4(s)

25. Practice Problem 5
Write balanced chemical equations for the following reaction: Solid sodium combines with chlorine gas to produce solid sodium chloride.
2Na(s) + Cl2(g) → 2NaCl(s)

26. Practice Problem 6
When solid copper reacts with aqueous silver nitrate, the products are aqueous copper(II) nitrate and solid silver.
Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)

27. Practice Problem 7
In a blast furnace, the reaction between solid iron(III) oxide and carbon monoxide gas produces solid iron and carbon dioxide gas.
Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g)

28. Types of Chemical Reactions
Section 2

29. 1. Decomposition Reactions
Decomposition reaction  a single compound has a reaction that makes two or more simpler substances
General equation AX  A + X
Most happen when energy (light/heat) is added

30. 2. Synthesis Reactions
Synthesis (composition) reaction  two or more substances combine to form a new compound
General equation A + X  AX

31. 3. Single-Replacement Reactions
Single-replacement reaction  one element replaces a similar element in a compound
General equation A + BX  AX + B
Y + BX  BY + X

32. 4. DOUBLE REPLACEMENT REACTIONS
Double-replacement reaction  ions of two compounds exchange places in an aqueous solution to form two new compounds
General equation AX + BY  AY + BX

33. 5. Combustion Reactions
Combustion reaction  substance combines with oxygen, releasing large amount of energy as heat and light
2H2(g) + O2(g) → 2H2O(g)

34. Reactions of Elements with Oxygen and Sulfur
One simple type of synthesis reaction is combination of element with oxygen to form an oxide of the element
Almost all metals react with oxygen to form oxides
Ex. Magnesium burned  magnesium oxide

35. 2Mg(s) + O2(g)  2MgO(s)

36. Group 2 Elements
Group 2 elements react to form oxides with general formula MO
M represents metal
Group 1 metals form oxides with general formula M2O
Li + O2  Li2O

37. Reactions with Sulfur
Groups 1 and 2 react with sulfur to make sulfides of the element
Group 1  M2S
Group 2  MS
16Rb(s) + S8(s) → 8Rb2S(s)
8Ba(s) + S8(s) → 8BaS(s)

38. Some metals (usually transition metals) combine with O2 to make two different oxides
Ex. Fe – can be +2 or +3
2Fe(s) + O2(g) → 2FeO(s)
4Fe(s) + 3O2(g) → 2Fe2O3(s)

39. Nonmetals Nonmetals also react with oxygen to make oxides
Sulfur reacts with oxygen to make sulfur dioxide
When carbon is burned, it makes carbon dioxide
S8(s) + 8O2(g) → 8SO2(g)
C(s) + O2(g) → CO2(g)

40. Hydrogen reacts with oxygen to make dihydrogen monoxide
2H2(g) + O2(g) → 2H2O(g)

41. Reactions of Metals & Halogens
Most metals react with halogens to make either ionic or covalent compounds
Ex. Group 1 reacts with halogens to form ionic compounds with formula MX
M = metal, X = halogen
2Na(s) + Cl2(g) → 2NaCl(s)
2K(s) + I2(g) → 2KI(s)
Reactions of Metals & Halogens

42. Group 2 Metals & Halogens
Formula MX2
Mg(s) + F2(g) → MgF2(s)
Sr(s) + Br2(l) → SrBr2(s)

43. Synthesis Reactions with Oxides
Active metals  highly reactive metals
Oxides of active metals react with water to make metal hydroxides
CaO(s) + H2O(l) → Ca(OH)2(s)
CaO = lime
Ca(OH)2 is important in setting cement
Synthesis Reactions with Oxides

44. Oxyacids Many oxides of nonmetals react with water to make oxyacids
SO2(g) + H2O(l) → H2SO3(aq)
In air polluted with SO2, reacts with oxygen in air to form sulfuric acid (acid rain)
2H2SO3(aq) + O2(g) → 2H2SO4(aq)

45. Practice Problems
Identify the products in each of the following reactions:
a. hydrogen burned in oxygen
H2O
b. H2(g) + N2(g) 
NH3
c. CaO(s) + H2O(l) 
Ca(OH)2(aq)

46. Decomposition of Binary Compounds
Simplest kind of decomposition reaction is binary compound into its elements
Ex. Passing electricity through water
2H2O(l) H2(g) + O2(g)
electricity

47. Electrolysis  decomposition of a substance by electricity

48. Oxides of less-active metals (lower center of PT) decompose into elements when heatedΔ
2HgO(s)  2Hg(l) + O2(g)

49. Decomposition of Metal Carbonates
When heated, metal carbonates break down to make a metal oxide and CO2 Δ
CaCO3(s)  CaO(s) + CO2(g)

50. Decomposition of Metal Hydroxides
All except with Group 1 metals decompose when heated to make metal oxides and water Δ
Ca(OH)2(s)  CaO(s) + H2O(g)

51. Decomposition of Metal Chlorates
When heated, metal chlorates decompose to make metal chloride and oxygen Δ
2KClO3(s) 2KCl(s) + 3O2(g)
MnO2(s)

52. Decomposition of Acids
Certain acids decompose into nonmetal oxides and water
H2CO3(aq) → CO2(g) + H2O(l) Δ
H2SO4(aq)  SO3(g) + H2O(l)

53. Practice Problems
Predict the products for these decomposition reactions
a. sodium chlorate 
Sodium chloride + oxygen
b. calcium carbonate 
Calcium oxide + carbon dioxide
c. potassium bromide 
Potassium + bromine

54. Replacement of a Metal in a Compound by Another Metal
Aluminum is more active than lead
When solid aluminum is placed in aqueous lead(II) nitrate, the aluminum replaces the lead
2Al(s) + 3Pb(NO3)2(aq) → 3Pb(s) + 2Al(NO3)3(aq)
Based on activity series of metals

56. Replacement of Hydrogen in Water by Metal
Most-active metals (Group 1) react strongly with water to make metal hydroxides and hydrogen
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Less-active metals react with steam or other form of energy

57. Replacement of Hydrogen in Acid by Metal
More-active metals react with certain acidic solutions and replace hydrogen
Reaction products are metal compound (salt) and hydrogen gas
Mg(s) + 2HCl(aq) → H2(g) + MgCl2(aq)

58. Replacement of Halogens
One halogen replaces another
Fluorine is most reactive
Can replace any other halogen
Cl2(g) + 2KBr(aq) → 2KCl(aq) + Br2(l)
F2(g) + 2NaCl(aq) → 2NaF(aq) + Cl2(g)
Br2(l) + KCl(aq) → no reaction

59. Practice Problems
For the following equations, predict what the products will be:
a. Ag + KNO3 →
No reaction
b. Zn + AgNO3 →
Zn(NO3)2 + Ag
c. Cl2 + KI →
I2 + 2KCl

60. d. Cu + FeSO4 →
No reaction
e. Fe + Pb(NO3)2 →
Pb + Fe(NO3)2
f. Cu + Al2(SO4)3 →
g. Al + Pb(NO3)2 →
Pb + Al(NO3)3
h. Cl2 + NaI →
I2 + NaCl

61. i. Fe + AgC2H3O2 →
Fe(C2H3O2)2 + Ag
j. Al + CuCl2 →
Cu + Al2Cl3
k. Br2 + CaI2 →
I2 + CaBr2
l. Fe + CuSO4 →
FeSO4 + Cu
m. Cl2 + MgI2 →
I2 + MgCl2

62. Formation of a Precipitate
Occurs when cations of one reactant combine with anions of another to form insoluble (or slightly soluble) compound
2KI(aq) + Pb(NO3)2(aq) → PbI2(s) + 2KNO3(aq)

63. Formation of a Gas
In some D-R reactions, one product is insoluble gas that bubbles out of mixture
FeS(s) + 2HCl(aq) → H2S(g) + FeCl2(aq)

64. Formation of Water In some D-R reactions, water is one of the products
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

65. Practice Problems
Classify each of the following reactions as synthesis, decomposition, single-replacement, double-replacement, or combustion:
a. N2(g) + 3H2(g) → 2NH3(g)
synthesis
b. 2Li(s) + 2H2O(l ) → 2LiOH(aq) + H2(g)
single-replacement
c. 2NaNO3(s) → 2NaNO2(s) + O2(g)
decomposition

66. d. 2C6H14(l ) + 19O2(g) →12CO2(g) + 14H2O(l )
combustion
e. NH4Cl(s) → NH3(g) + HCl(g)
decomposition
f. BaO(s) + H2O(l ) → Ba(OH)2(aq)
synthesis
g. AgNO3(aq) + NaCl(aq) →AgCl(s) + NaNO3(aq)
double-replacement

67. Practice Problem
For each of the following reactions, identify the missing substances, then balance the final equation. Each slot may be one OR MORE substances.
a. synthesis: _____ → Li2O
4Li + O2  2Li2O
b. decomposition: Mg(ClO3)2 → _____
Mg(ClO3)2  MgCl2 + 3O2

68. c. single-replacement: Na + H2O → _____
2Na + 2H2O  2NaOH + H2
d. double-replacement: HNO3 + Ca(OH)2 → _____
2HNO3 + Ca(OH)2  Ca(NO3)2 + 2H2O
e. combustion: C5H12 + O2 → _____
C5H12 + 8O2  5CO2 + 6H2O