1. Chapter 11
Chemical Reactions

2. Chemical Equation Describes chemical reaction.
Chemical equation: reactants yield products
Reactants  Products
Much easier to write symbols and formulas instead of words

3. Examples CH4(g) + O2(g)  CO2(↑) + H2O(↑)
Solid Iron reacts with oxygen gas to form the solid IronIIIoxide.
iron(s) + oxygen(g)  ironIIIoxide(s)
Fe(s) + O2(g)  Fe2O3(s)
Carbon tetrahydride gas BURNS to form carbon dioxide gas and water vapor.
Carbon tetrahydride(g) + oxygen(g)  carbon dioxide(↑) + water(↑)
CH4(g) + O2(g)  CO2(↑) + H2O(↑)
Skeleton Equation: chemical equation that tells you what the reactants and products are but NOT how much of each you have.
First step in writing a chemical equation.

4. Symbols Used (s) solid (l) liquid (g) gas (↑) gas as a product
(aq) aqueous (in water solution)
() ppt (precipitate) solid product from 2 aqueous reactants
D means with heat
Pt means with Platinum catalyst: speeds up a reaction without being used.
 reversible reaction

5. Balancing Chemical Equations
Balanced equations have:
the same # of atoms of each element on BOTH sides of the equation.
Law of Conservation of Mass – atoms can neither be created nor destroyed, simply rearranged.

6. Rules for Balancing Equations
Get the correct formulas for reactants and products.
(USE ION CHART AND DON”T FORGET DIATOMIC ELEMENTS!)
Write reactants on left, products on right. Use plus signs to separate compounds and yield sign to separate the reactants from products.

7. Rules Continued
Count the # of atoms of each element in reactants and products. (Polyatomic atoms on both sides count as one.)
Balance # of each element using coefficients.
Coefficient – small whole # in front of a formula.
NEVER CHANGE FORMULA SUBSCRIPTS

8. Rules for Balancing Equations
Balance elements appearing 3 or more places LAST.
Check each element to make sure equation is balanced.
Make sure all coefficients are in the lowest whole number ratio.
Do not change subscripts!!!

9. H O N Cl Br I F Diatomic Molecules There are 7 naturally occurring
Diatomic Molecules- a molecule made up two atoms of the same element. They are only diatomic when they are alone.
There are 7 naturally occurring
diatomic molecules.
H O N Cl Br I F

10. Balancing Examples ___ C(s) + ___ O2(g)  ___ CO2 (g)
___ AgNO3(aq) + ___Cu(s)  ___ Cu(NO3)2(aq) + ___ Ag(s)
___ Al(s) + ___ O3(g)  ___ Al2O3(s)
*___ C2H6(g) + ___ O2(g)  ___ CO2(g) + ___ H2O(g)
*___ H3PO3  ___ H3PO4 + PH3

11. 5 Types of Chemical Reactions
Combination Reaction – elements combine to form a compound.
A + B AB
element + element  compound
Ex. Sodium + chlorine  sodium chloride
___Na(s) + ___ Cl2(g)  ___ NaCl(s) 2 2

12. 5 Types of Reactions AB  A + B compound  element + element
Decomposition Reaction – compound breaks down into its element.
AB  A + B
compound  element + element
Ex: MercuryII oxide  mercury + oxygen
___ HgO  ___Hg + ___O2 2 2

13. AB + C A + CB or AB + D AD + B 5 Types of Reactions - 3
Single Replacement Reaction – one element replaces another element in a compound.
AB + C A + CB or
AB + D AD + B

14. Examples of Single Replacement Reactions
Must use Activity Series to see if reaction works
Zinc + sulfuric acid  zinc sulfate + hydorgen
Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2(↑)
Periodic table is activity series for halogens
Sodium bromide + chlorine  sodium chloride + bromine
___NaBr(s) + ___Cl2(g)  ___NaCl(s) + ___Br2(↑) 2 2

15. AB + CD AD + CB 5 Types of Reactions
Double Replacement Reaction – two compounds react and exchange positive ions to form two new compounds.
AB + CD AD + CB
Barium Chloride(aq) + potassium carbonate(aq)  barium carbonate() + potassium chloride(aq)
BaCl2(aq) + K2CO3(aq)  BaCO3() + ___ KCl(aq) 2

16. CxHy + O2 CO2 + H20 5 Types of Reactions
Combustion Reaction – oxygen reacts with a compound composed of C and H.
CxHy + O CO2 + H20
Also called Burning (exothermic)
The products are always CO2 and H2O.

17. Examples of Combustion Reactions
C6H O CO H2O 7½ 6 3
C6H O CO H2O
CH3OH O CO H2O 1½ 2
CH3OH O CO H2O

18. Special Decomposition Reactions:
Decomposition of a Carbonate:
Metal carbonate  metal oxide + carbon dioxide
XCO XO + CO2
ex. Na2CO Na2O + CO2

19. Special Decomposition Reactions:
Decomposition of a Hydroxide:
Metal hydroxide  metal oxide + water
XOH XO + H2O
ex. 2NaOH Na2O + H2O

20. Special Decomposition Reactions:
Decomposition of a Chlorate: (ClO3)
Metal chlorate  metal chloride + oxygen
XClO XCl + O2
ex ___NaClO ___NaCl + ___O2 2 3

21. Special Decomposition Reactions: 4
Special single Replacement Reaction:
Group IA or IIA metal and H2O
X + HOH XOH + H2
ex. 2Na + 2HOH NaOH + H2

22. How to ID types of reactions.
Combination Reactions – given 2 items that form 1 new compound.
Decomposition Reactions – given a single compound that breaks into parts.
Single Replacement – given a single element plus a single compound, forms a new compound a a different element.
Double Replacement – given two compounds (+’s change places).
Combustion Reaction – given CH compound with Oxygen, always forms water and carbon dioxide.

23. The End