1. Chemical reactions
Chapter 11

2. Section Overview 11.1: Describing Chemical Reactions
11.2: Types of Chemical Reactions
11.3: Reactions in Aqueous Solution

3. Describing Chemical reactions
Section 11.1

4. Describing Chemical Reactions
Chemical reactions are necessary to sustain life.
In a chemical reaction, one or many substances change (reactants) into one or more new substances (products).
Chemists use a chemical equation to convey as much as possible about what happens in a chemical reaction.
Word Equations: Reactants are on the left and products on the right with an arrow between them that means yields, or gives, or reacts to produce.
Reactants  Products
Iron + oxygen  iron (III) oxide
Hydrogen peroxide  water + oxygen

5. Describing Chemical Reactions
Chemical Equations: A representation of a chemical reaction using the formulas for the reactants and products instead of words.
Fe + O2  Fe2O2
Equations that just show the formulas and not the amount of each of the reactants and products are called skeleton equations.
To add more information, you can also add the state of the of substances in parentheses.
Fe(s) + O2(g)  Fe2O2(s)
Most chemical reactions require a catalyst, which speeds up the reaction and is written above the arrow.

6. Writing a Skeleton Formula
Example Problem: Hydrochloric acid and solid sodium hydrogen carbonate are shown before being placed into a beaker to react. The products formed are aqueous sodium chloride, water, and carbon dioxide gas. What is the skeleton formula for this reaction?

7. Writing a Skeleton Equation
Example Problem: Hydrochloric acid and solid sodium hydrogen carbonate are shown before being placed into a beaker to react. The products formed are aqueous sodium chloride, water, and carbon dioxide gas. What is the skeleton formula for this reaction?
Solution: Write the formulas for everything given and put into a skeleton formula.
Hydrochloric acid: HCl(aq) Water: H2O(l)
Carbon dioxide gas: CO2(g) Sodium chloride: NaCl(aq)
Sodium hydrogen carbonate: NaHCO3(g)
NaHCO3(s) + HCl(aq)  NaCl(aq) + H2O(l) + CO2(g)

8. Balancing Chemical Equations
A balanced equation is a chemical equation in which both sides have equal numbers of atoms of each element.
To balance an equation, you place coefficients, small whole numbers that are placed in front of the formulas in an equation in order to balance it.
To write a balanced chemical equation, first write the skeleton equation.
Then use coefficients to balance the equation.
Occasionally, a skeleton equation may already be balanced.

9. Balancing Chemical Equations
Determine the correct formulas for all the reactants and products.
Write a skeleton equation.
Determine the number of atoms each side of the equation. Count a polyatomic ion as a single unit if it appears unchanged on both sides.
Balance the elements one at a time using coefficients. Begin with elements that only appear once on each side. Never balance by changing subscripts!
Check each atom or polyatomic ion to be sure they are balanced.
Make sure the coefficients are in the lowest possible ratio.

10. Balancing Chemical Equations
Example Problem 1: Hydrogen and oxygen react to form water. Write a balanced equation for the reaction.

11. Balancing Chemical Equations
Example Problem 1: Hydrogen and oxygen react to form water. Write a balanced equation for the reaction.
Solution: Skeleton Formula First
H2 + O2  H2O
H = 2 H=2
O = 2 O=1
Skeleton isn’t balanced, so start adding coefficients.
2H2 + O2  H2O
H= H=2
O= O=1
Still un balanced
2H2 + O2  2H2O
H= H=4
O= O=2

12. Balancing Chemical Equations
Example Problem 2: Balance the following equation
Al + O2  Al2O3

13. Balancing Chemical Equations
Example Problem 2: Balance the following equation
Al + O2  Al2O3
Solution:
Al + O2  Al2O3 4Al + O2  2Al2O3
Al =1 Al=2 Al= Al=4
O=2 O=3 O= O=6
2Al + O2  Al2O3 4Al + 3O2  2Al2O3
Al=2 Al=2 Al= Al=4
O=2 O=3 O= O=6
2Al + O2  2Al2O3
Al= Al=4
O= O=6

14. Types of Chemical reactions
Section 11.2

15. Classifying Chemical Reactions
The five types of chemical reactions are combination, decomposition, single-replacement, double-replacement, and combustion.
Combination Reactions: A chemical change in which one or more substances react to form a new single substance.
R + S  RS
2K + Cl2  2KCl
Decomposition Reaction: A chemical change in which a single compound breaks down into two or more simpler compounds.
RS  R + S
2HgO  2Hg + O2

16. Classifying Chemical Reactions
Single-Replacement Reaction: A chemical change in which one element replaces a second element in a compound.
T + RS  TS + R
2K + 2H2O  2KOH + H2
Whether one metal will replace another depends upon the reactivities of the two metals.
A reactive metal will replace any metal listed below it in an activity series.

17. Classifying Chemical Reactions

18. Classifying Chemical Reactions
Double Replacement Reaction: A chemical change involving an exchange of positive ions between two compounds.
R+S- + T+U-  R + U- + T+S-
Na2S + Cd(NO3)2  CdS + 2NaNO3
Combustion Reaction: A chemical change in which an element or a compound reacts with oxygen, often producing energy in the form of heat and light.
CXHY + (x+y/4)O2  xCO2 + (y/2)H2O
CH4 + 2O2  CO2  2H2O

19. Writing Equations for Combination Reactions
Example Problem: Complete the equation for the following reaction Cu + S  (two reactions possible)

20. Writing Equations for Combination Reactions
Example Problem: Complete the equation for the following reaction Cu + S  (two reactions possible)
Solution:
Two reactions are possible because copper is a transition metal and can be Cu+ or Cu2+.
Cu + S  CuS (balanced) OR
Cu + S  Cu2S
2Cu + S  Cu2S (balanced)

21. Writing Equations for Decomposition Reactions
Example Problem: Write a balanced chemical equation for the following decomposition reaction H2O 

22. Writing Equations for Decomposition Reactions
Example Problem: Write a balanced chemical equation for the following decomposition reaction H2O 
Solution:
H2O  H2 + O2
H= H=2
O= O=2
2H2O  H2 + O2
H= H=2
O= O=2
2H2O  2H2 + O2 (balanced)

23. Writing Equations for Single Replacement Reactions
Example Problem: Write a balanced equation for each reaction.
Zn + H2SO4 
Cl2 + NaBr 

24. Writing Equations for Single Replacement Reactions
Example Problem: Write a balanced equation for each reaction.
Zn + H2SO4 
Cl2 + NaBr 
Solution:
Zn + H2SO4  ZnSO4 + H2 (balanced)
Zinc is more reactive than hydrogen, so it takes its place.
b. Cl2 + NaBr  NaCl + Br2
Chlorine is more reactive than bromine (reactivity decreases down a group).
Cl2 + 2NaBr  2NaCl + Br2 (balanced)

25. Writing Equations for Double-Replacement Reactions
Example Problem: Write a balanced chemical equation for each double replacement reaction
CaBr2 + AgNO3  (a precipitate of silver bromide is formed)
FeS + HCl  (hydrogen sulfide gas (H2S) is formed)

26. Writing Equations for Double-Replacement Reactions
Example Problem: Write a balanced chemical equation for each double replacement reaction
CaBr2 + AgNO3  (a precipitate of silver bromide is formed)
FeS + HCl  (hydrogen sulfide gas (H2S) is formed)
Solution:
CaBr2 + AgNO3  AgBr + Ca(NO3)2
CaBr2 + 2AgNO3  2AgBr + Ca(NO3)2 (balanced)
FeS + HCl  H2S + FeCl2
FeS + 2HCl  H2S + FeCl2 (balanced)

27. Writing Equations for Combustion Reactions
Example Problem: Write balanced equations for the complete combustion of these compounds
Benzene (C6H6)
Ethanol (CH3CH2OH)

28. Writing Equations for Combustion Reactions
Example Problem: Write balanced equations for the complete combustion of these compounds
Benzene (C6H6)
Ethanol (CH3CH2OH)
Solution: Remember, oxygen is the other reactant. Products are CO2 and H2O.
C6H6 + O2  CO2 + H2O
2C6H6 + 15O2  12CO2 + 6H2O (balanced)
CH3CH2OH + O2  CO2 + H2O
CH3CH2OH + 3O2  2CO2 + 3H2O (balanced)

29. Reactions in aqueous soultion
Section 11.3

30. Net Ionic Equations
Most chemical reactions take place in water, in which some of the reactants dissociate, or separate, into cations and anions when they dissolve in water.
These ions can be used to write a complete ionic equation, an equation that shows dissolved ionic compounds as dissociated free ions.
An ion that appears on both sides of the equation and is not directly involved in the reaction is called a spectator ion.
When you rewrite the equation without the spectator ions, the net ionic equation remains.

31. Net Ionic Equations AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
Complete Ionic Equation
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) 
AgSl(s) + Na+(aq) + NO3-(aq)
The nitrate and sodium ions remain unchanged on both sides and are spectator ions therefore they can be removed.
Net Ionic Equation
Ag+(aq) + Cl-(aq)  AgSl(s)
Be sure in the balanced equation that not only the atoms are equal, but also the charges.

32. Net Ionic Equations
Example Problem: Aqueous solutions of iron (III) chloride and potassium hydroxide are mixed. A precipice of iron (III) hydroxide forms. Write the balanced net ionic equation for the reaction.

33. Net Ionic Equations
Example Problem: Aqueous solutions of iron (III) chloride and potassium hydroxide are mixed. A precipice of iron (III) hydroxide forms. Write the balanced net ionic equation for the reaction.
Solution:
Fe3+(aq) + Cl- (aq) + K+(aq) + OH-(aq) 
Fe(OH)3(s) + K+(aq) + Cl-(aq)
Spectator ions = Cl and K
Fe3+(aq) + OH-(aq)  Fe(OH)3(s)
Fe3+(aq) + 3OH-(aq)  Fe(OH)3(s) (balanced)

34. Predicting the Formation of a Precipitate
Sometimes mixing solutions of two ionic compounds can result in the formation of an insoluble salt called a precipitate.
Whether or not a precipitate forms, depends on the solubility of the new compounds that form.
You can predict the formation of a precipitate by sing the general rules for solubility of ionic compounds.
If a compound is soluble, it will not form a precipitate.
If a compound is insoluble, it will form a precipitate.