1. Balancing Chemical Equations

2. Balanced Equations Atoms can not be created or destroyed
All atoms we start with we must end up with…and vice versa!
A balanced equation has the same number of each element on both sides of the equation.

3. But what if an equation isn’t already balanced?
C + O2  CO2
This equation is already balanced
But what if an equation isn’t already balanced?

4. We need one more ________ in the products.
Like this….
C + O2  CO
We need one more ________ in the products.
You can’t change the formula, because it describes what is being produced…CO (Carbon Monoxide.)

5. The other oxygen must be used to make another CO
This gives a better/closer idea of what is happening…
BUT….
The other oxygen must be used to make another CO
But where does the other C come from?

6. Must have started with two C’s
2 C + O2  2 CO

7. Rules for Balancing
Write the correct formulas for all the reactants and products
Count the number of atoms of each type appearing on both sides.
Balance the elements one at a time by adding coefficients (the numbers in front)
2 CO2
Check to see if it is balanced

8. Never Never change a subscript to balance an equation CO2
If you change the formula you are describing a different reaction.
H2O is a different compound than H2O2
Never put a coefficient in the middle of a formula
2 NaCl is ok Na2Cl is not.

9. Example
H2 + O2  H2O
Make a table to keep track of atoms

10. Example
H2 + O2  H2O R P 2 H O 1
Need twice as much O in the product

11. Example
H2 + O2  2H2O R P 2 H O 1
Changes the O

12. Example
H2 + O2  2H2O R P 2 H O 1 2
Also changes the H

13. Example
H2 + O2  2H2O R P 2 H O 1 4 2
Now we need twice as much H in the reactant

14. Example
2H2 + O2  2H2O R P 2 H O 1 4 2
Recount to check

15. Example
2H2 + O2  2H2O
Your answer R P 2 H O 1 4 4 2
Recount to check

16. Types of Reactions Millions of reactions Too many to remember
They fall into several categories
We will focus on Double Replacement in today’s lab

17. Double Replacement Two things replace each other
Reactants must be two ionic compounds or acids.
Usually in aqueous solution
NaOH + FeCl3 
The positive ions change place
NaOH + FeCl3  Fe+3OH- +Na+1Cl-1
NaOH + FeCl3  Fe(OH)3 + NaCl

18. Double Replacement Will only happen if one of the products
Doesn’t dissolve in water and forms a solid
(look at solubility rules)
Or is a gas that bubbles out
Or is a covalent compound usually water
After adding lead nitrate
Potassium iodide
2KI(aq) + Pb(NO3)2 (aq)  2KNO3(aq) + PbI2 (s)
PbI2 lead (II) iodide is insoluble

19. General Rules for the Water Solubilities of Common Ionic Compounds
Compounds that are mostly soluble:
All nitrates
Alkali metal (group 1A) and ammonium compounds
Chlorides, bromides, and iodides, except for those of Pb2+, Ag+, Hg2+
Sulfates except for those of Sr2+, Ba2+, Pb2+, and Hg2+
CaSO4 is slightly soluble

20. General Solubility Rules
Compounds that are mostly insoluble:
Carbonates, hydroxides, and sulfides, except for ammonium compounds and those of the group 1A metals. (The hydroxides and sulfides of Ca2+, Sr2+, and Ba2+ are slightly to moderately soluble.)