Chemistry 10050 - Chemical Reactions
Definitions
- Reactant:
- A starting material for a chemical reaction. Consumed during course of
reaction.
- Product:
- A substance formed by a chemical reaction.
- Avogadro's Number:
- ≈6.02 x 1023.
- Mole:
- The amount of a substance that contains Avogadro's number of
"units".
- Molecular Weight:
- The sum of atomic weights in a molecule.
- Formula Weight:
- The sum of atomic weights in a formula unit. (For ionic compounds).
- Percent Yield:
- Measure of amount of compound obtained in lab. (Given as a percentage)
Chemical Reaction Example
| Reactants |
Reaction |
Products |
| 2 CH3CO2H(l) +
Na2CO3(s) |
→ |
CO2(g) +
H2O(l) +
2 NaCH3CO2(aq) |
Coefficients - indicate the number of each type of molecule/ion
required for complete reaction.
In this example, two molecules of CH3CO2H react with one 'molecule'
of Na2CO3.
| Showing Physical States |
| (s) |
(l) |
(g) |
(aq) |
| solid |
liquid |
gas |
aqueous solution |
Balancing Chemical Reactions
- Write unbalanced equation, including all compounds with correct formulas. The chemical
formulas cannot be changed to balance the equation.
- Select "easiest" element, and balance the number of this
atom in reactants with the number of this atom in products
by changing the coefficients.
- Repeat step (2) for each of the remaining elements. Verify that the simplest whole
number ratio found.
- Check results.
Mole Relationships
A mole is a number, not a weight. Just as a ream of paper comes in different weights, a
mole of hydrogen does not weigh the same as a mole of oxygen. Chemical reactions are best
thought of as the interactions of single molecules. To convert to a scale that is
reasonable for actually performing the experiment, we typically use moles of each
compound, which have the same relationships to chemical equations as the numbers of
compounds used in the balanced chemical equation.
To convert between moles of a compound and weight in grams, use the molecular weight
(in grams/mole). To convert between moles of one substance and moles of another, use the
coefficients of the balanced equation.
|
a A + b B + ... |
→ |
c C + d D + ... |
|
|
mass (g) |
|
mass (g) |
|
| (from molecular weight in g/mole) |
¯ |
|
|
(from molecular weight in g/mole) |
|
moles |
→ |
moles |
|
|
(from coefficients of balanced equation) |
|
Example: How many grams of H2O can be produced from
3.75 g of O2?
| M.W. |
32.0 g/mol |
|
2.02 g/mole |
|
18.0 g/mol |
|
|
O2 |
+ |
2 H2 |
|
2 H2O |
|
|
3.75 g |
|
|
|
4.22 g |
|
| 3.75 g x (1 mole O2)/(32.0 g) |
↓ |
|
|
|
↑ |
0.234 moles H2O x (18.0 g H2O)/(1 mole
H2O) |
|
0.117 moles O2 |
|
→ |
|
0.234 moles H2O |
|
|
0.117 mole O2 x (2 mole H2O)/(1
mole O2) |
|
Classification of Chemical Reactions:
- Precipitation: reaction of two ionic compounds to produce an insoluble ionic
salt.
| Solubility Rules: |
| NH4+, Li+, K+, Na+,
NO3- |
salts always soluble. |
| Cl-, Br-, I- |
salts soluble (except with Ag+, Hg22+,
and Pb2+) |
| SO42- |
salts soluble (except with Ag+, Hg22+,
and Pb2+) |
| CO32-, PO43-, S2- |
salts are always insoluble (unless one of above rules applies) |
| OH- |
salts are always insoluble (unless with Ba2+ or one of above
rules applies) |
- Acid-Base: reaction of an acid with a base to yield a salt plus water.
- Acids: HCl, HBr, HI, HNO3, and H2SO4
Bases: Any cation combined with OH- (example: NaOH, CaOH, Ba(OH)2,
etc.)
- Oxidation-Reduction: reaction in which electrons are transfered between
molecules.
- Oxidation: Loss of one or more electrons.
Reduction: Gain of one or more electrons.
Oxidizing Agent: Causes something else to be/b> Causes something else to be oxidized (it gets reduced).
Reducing Agent: Causes something else to be reduced (it gets oxidized).
Oxidation Number: Formal system for assigning location of electrons.
Ionic Equations:
- Molecular Equation: All compounds (including ionic salts) written as molecules.
- Ionic Equation: All salts written with ions separated (as actually found in
solution).
- Net Ionic Equation: Ionic equation with ions that do not undergo any reaction
removed.
- Example: (Same reaction, with molecular, ionic, and net ionic equations given)
Pb(NO3)2(aq) + 2 NaCl(aq)
→ PbCl2(s) + 2 NaNO3(aq)
Pb+2(aq) + 2 NO3-(aq) + 2 Na+(aq)
+ 2 Cl-(aq) → PbCl2(s)
+ 2 Na+(aq) + 2 NO3-(aq)
Pb+2(aq) + 2 Cl-(aq)
→ PbCl2(s)
Rules for Assigning Oxidation States:
- All atoms in an element have an oxidation state of zero (0).
- A single atom ion has an oxidation state equal to its charge.
- Most electronegative atom(s) have oxidation state equal to that expected if ionic.
(Ex. F-1, Cl-1, Br-1, I-1,
O2-)
- Sum of oxidation state must equal total charge.
- Examples:
- Na: Na = Na0
- SO3: S = S+6, O = O-2.
- Cu(NO3)2 = (Cu+2)(NO3-)2:
Cu = Cu+2, N = N+5, O = O-2.