1. NOTES: 11.1 – Writing and Balancing Chemical Equations
Chemical Reactions

2. Effects of chemical reactions:
• Chemical reactions rearrange atoms in the reactants to form new products.
• The identities and properties of the products are completely different from that of the reactants.

3. What is a Chemical Reaction?
• Chemical Reaction – one or more substances change into new substances
• Process involves reactants and products
• Reactant – a starting substance
• Product – a substance formed
• Example:
Nitrogen and hydrogen gas can react to form ammonia under certain conditions.
Reactants Yield Products
N2 (g) H2 (g)  NH3 (g)

4. How Can You Tell Whether or Not a Chemical Reaction Has Taken Place?
• Chemical Change – alters a given material by changing its chemical composition
• Production of gases and color changes are signs of chemical reactions
• Examples: burn, rust, decompose, corrode, explode

5. Chemical Reactions: • we can describe a chemical reaction with words:
“Iron metal reacts with oxygen gas to produce iron(III) oxide, or rust.”
• we can then write a word equation:
Iron + oxygen  iron(III) oxide

6. Chemical Equations: 2H2 + O2  2H2O
• Chemical equations are used to represent or describe chemical reactions.
• Chemical equations use chemical symbols and formulas for the reactants and products
• For example when hydrogen, H2,burns, it reacts with oxygen, O2, in the air to form water. We write the chemical equation for this reaction as follows:
2H2 + O2  2H2O

7. Chemical Equations: An equation shows…
 Chemical formulas of reactants;
 Chemical formulas of products;
 Molecule / Mole ratios of all compounds in the reaction.

8. Chemical Equations:
We read the (+) sign as “reacts with” and the arrow ( ) as “produces” or “yields”.
2H2 + O2  2H2O
Reactants
Products

9. To show physical states of each substance:
• (s) = solid
• (l) = liquid
• (g) = gas
• (aq) = aqueous
**aqueous means dissolved in water

10. • Consider again the reaction of iron with oxygen to form iron(III) oxide, or rust.
Fe(s) + O2 (g)  Fe2O3 (s)
(unbalanced)
**this is a skeleton equation in that is NOT “balanced” and does not show the relative amounts of reactants and products

11. Coefficients & Subscripts
COEFFICIENTS: numbers in front of compound that represents the number of molecules/moles of that compound
SUBSCRIPTS: small numbers within a formula that help define the compound.
2H2SO4
Subscript
Coefficient

12. H2O: One molecule of water
2H2O: Two molecules of water
H2O2: One molecule of Hydrogen Peroxide

13. • During a chem. rxn. atoms are rearranged (NOT created or destroyed!)
• Chemical equations must be BALANCED to show the relative amounts of all substances.
• Balanced means: each side of the equation has the same # of atoms of each element.
CH4 + O2  H2O + CO2 Unbalanced
CH4 + 2O2  2H2O + CO2 Balanced

14. In order to balance…
• Write correct formulas for all reactants and products
• Reactants  Products
• Count the number of atoms of each element in reactants & products.
• Balance one at a time using coefficients.
• Check for balance
• Are the coefficients in the lowest possible ratio?

15. Balancing Equations
NOTE: When balancing equations, you may change coefficients as much as you need to, but you may never change subscripts because you can’t change what substances are involved.

16. Fe(s) + O2(g)  Fe2O3(s) (unbalanced)
4Fe(s) + 3O2(g)  2Fe2O3(s) (balanced)

17. Sample Problem 1:
● Water is decomposed (broken down) to form the gaseous products hydrogen, H2, and oxygen, O2. Write the balanced equation for this reaction.
● H2O  H2 + O2
**O is not balanced
● 2H2O  2H2 + O2
**The equation is balanced!

18. Sample Problem 2:
● Chlorine gas, Cl2, reacts with potassium bromide, KBr, to form potassium chloride and bromine, Br2. Write the balanced equation for this reaction,
● Cl2 + KBr  KCl + Br2
**Cl and Br are not balanced
● Cl2 + 2KBr  2KCl + Br2
**The equation is balanced!

19. Examples: (3:2:3:2) CuCl2 (aq) + Al (s)  Cu (s) +AlCl3 (aq)

20. Balance C – then H – then O
Examples:
Propane, C3H8, burns in oxygen, O2, to form carbon dioxide and water.
C3H8 + O2  CO2 + H2O
Balance C – then H – then O
C3H8 + 5O2  3CO2 + 4H2O
(1:5:3:4)

21. Balance C – then H – then O
Examples:
Pentane, C5H12, burns in oxygen, O2, to form carbon dioxide and water.
C5H12 + O2  CO2 + H2O
Balance C – then H – then O
C5H12 + 8O2  5CO2 + 6H2O
(1:8:5:6)

22. Examples:
Silver nitrate reacts with copper to produce silver and copper (II) nitrate.
AgNO3 + Cu  Ag + Cu(NO3)2
2AgNO3 + Cu  2Ag + Cu(NO3)2
(2:1:2:1)

23. Examples:
Phosphorus reacts with oxygen gas to produce diphosphorus pentoxide.
P + O2  P2O5
4P + 5O2  2P2O5
(4:5:2)

24. Balance C – then H – then O
Examples:
C7H14 + O2  CO2 + H2O
Balance C – then H – then O
C7H ½O2  7CO2 + 7H2O
2C7H O2  14CO2 + 14H2O
(2:21:14:14)