1. Chemical Reactions reactants products
Balancing chemical equations
Types of chemical reactions

2. How reactants are transformed into products?
Reactants are transformed during chemical reactions
Energy is required (absorbed) to break a chemical bond
Energy is released when a chemical bond forms

3. Balancing equations Obey law of conservation of matter
Chemical equations have two parts: reactants and products

4. How to balance?
The total number of atoms of each element should be the same on both sides of equation
Use coefficients to balance equations
Example:
2H2 + O2 2 H2O

5. Types of chemical reactions
Single replacement
Double replacement
Decomposition
Combination
Combustion

6. Synthesis (Combination) Reactions
Two or more substances combine to form a
new compound.
A + X  AX
Reaction of elements with oxygen and sulfur
Reactions of metals with Halogens
Synthesis Reactions with Oxides
There are others not covered here!

7. Decomposition Reactions
A single compound undergoes a reaction that
produces two or more simpler substances
AX  A + X
Decomposition of:
  Binary compounds 2H2O(l )  2H2(g) + O2(g)
  Metal carbonates CaCO3(s)  CaO(s) + CO2(g)
  Metal hydroxides Ca(OH)2(s)  CaO(s) + H2O(g)
  Metal chlorates 2KClO3(s)  2KCl(s) + 3O2(g)
  Oxyacids H2CO3(aq)  CO2(g) + H2O(l )

8. Decomposition Reactions
Sulfates
With the exception of alkali metals and alkaline sulfates, all other metals are decomposed by heat to form a metal oxide
Nitrates
Alkali metals decompose on heating to yield the nitrites and oxygen. All other metal nitrates are decomposed to nitrogen dioxide, oxygen, and the metal oxide on heating.

9. Single Replacement Reactions
A + BX  AX + B
BX + Y  BY + X
Replacement of:
Metals by another metal
Hydrogen in water by a metal
Hydrogen in an acid by a metal
Halogens by more active halogens

10. The Activity Series of the Metals
  Lithium
  Potassium
  Calcium
  Sodium
  Magnesium
  Aluminum
  Zinc
  Chromium
  Iron
  Nickel
  Lead
  Hydrogen
  Bismuth
  Copper
  Mercury
  Silver
  Platinum
  Gold
Metals can replace other metals
provided that they are above the
metal that they are trying to
replace.
Metals above hydrogen can
replace hydrogen in acids.
Metals from sodium upward can
replace hydrogen in water

11. Predict if these reactions will occur3
Mg AlCl3 2 2
Al MgCl2 3
Can magnesium replace aluminum?
YES, magnesium is more reactive than aluminum.
Activity Series
Al MgCl2
No reaction
Can aluminum replace magnesium?
NO, aluminum is less reactive than magnesium.
Therefore, no reaction will occur.
Activity Series
Order of reactants
DOES NOT
determine how
they react.
MgCl Al
No reaction
The question we must ask is can the single element replace its counterpart?
metal replaces metal or nonmetal replaces nonmetal.

12. The Activity Series of the Halogens
  Fluorine
  Chlorine
  Bromine
  Iodine
Halogens can replace other
halogens in compounds, provided
that they are above the halogen
that they are trying to replace.
2NaCl(s) + F2(g) 
2NaF(s) + Cl2(g)
???
MgCl2(s) + Br2(g) 
???
No Reaction

13. Single Replacement Reactions
Single Replacement Reactions
Sodium chloride solid reacts with fluorine gas
NaCl(s) + F2(g)  NaF(s) + Cl2(g)
Note that fluorine replaces chlorine in the compound
Aluminum metal reacts with aqueous copper (II) nitrate
Al(s)+ Cu(NO3)2(aq) Cu(s) + Al(NO3)3(aq) 2 2 2 3 3 2

14. Double Replacement Reactions
The ions of two compounds exchange places in an
aqueous solution to form two new compounds.
AX + BY  AY + BX
One of the compounds formed is usually a
precipitate, an insoluble gas that bubbles out of
solution, or a molecular compound, usually water.

15. Double Replacement Reactions
Double Replacement Reactions
Think about it like “foil”ing in algebra, first and last ions go together + inside ions go together
Example:
AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq)
Another example:
K2SO4(aq) + Ba(NO3)2(aq)  KNO3(aq) + BaSO4(s) 2

16. Combustion Reactions
A substance combines with oxygen, releasing a large
amount of energy in the form of light and heat.
  Reactive elements combine with oxygen
P4(s) + 5O2(g)  P4O10(s)
(This is also a synthesis reaction)
  The burning of natural gas, wood, gasoline
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)

17. Write a balanced chemical equation for the following combustion reactions:
A. C5H12
O2  CO H2O 8 5 6
B. C22H46 2
O2  CO H2O 67 44 46
C. C15H28
O2  CO H2O 22 15 14

18. Solubility and precipitation reactions
Pb(NO3)2 (aq) + 2NaI(aq)  PbI2(s) NaNO3(aq)

19. Predict if a reaction will occur when you combine aqueous solutions
of iron (II) chloride with aqueous sodium carbonate solution.
If the reaction does occur, write a balanced chemical equation showing it.
(be sure to include phase notation)
iron (II) chloride sodium carbonate
sodium chloride +
iron (II) carbonate
Fe2+
Cl1-
Na1+
CO32-
Fe2+
Cl1-
Na1+
CO32-
Cl2 Fe
CO3
Na2
NaCl
FeCO3
(aq)
(ppt)
Using a SOLUBILITY TABLE:
sodium chloride is soluble
iron (II) carbonate is insoluble
Balanced chemical equation
FeCl2
Na2CO3
NaCl
FeCO3
(aq)
(ppt) + 2
(aq)
(aq)
Complete Ionic Equation
Fe2+(aq) + 2Cl1-(aq) + 2Na1+(aq) + CO32-(aq) Na1+(aq) + 2Cl1-(aq) + FeCO3(s)

20. Solubility rules Soluble in water:
sodium, potassium, and ammonium salts; acetates and nitrates
Halides with the exception of halides of lead (II), silver(I), and mercury(I).
Sulfates with the exception of sulfates of calcium, barium, lead (II) and strontium

21. Insoluble in water
Phosphates, carbonates and sulfides except sodium, potassium, ammonium salts, and calcium sulfide
Hydroxides except sodium, potassium, calcium, and barium hydroxides

22. Acid-Base reactions (neutralization reactions)
Acid: any compound that produces hydrogen ions (H+), when added to water.
Base: any substance that produces hydroxide ions (OH-), when added to water.
HCl(aq) + Na(OH)(aq)  NaCl(aq) + H2O(l)