CHEM 105N - Intro to Chemical Reactions Notes

Chemical Reactions

Law of Conservation of Matter

• Matter cannot be created nor destroyed.

• The amount of matter in a system stays constant.

Chemical Equations

• Represent chemical reactions on paper.

• Reactants (starting materials) are on the left, products (ending materials) are on the right, separated by an arrow.

• "+" separates multiple reactants or products.

• Coefficients are added to balance equations, not subscripts, to avoid changing the compound.

Balancing Equations

• Follow the Law of Conservation of Mass.

• Start with an element in only one reactant and product.

• Balance by changing coefficients, not subscripts.

• Move to other elements until complete, checking all elements at the end.

States of Matter

• States of matter are often written in parentheses to the right of each formula or symbol.

  • (g) = gas

  • (l) = liquid

  • (s) = solid

  • (aq) = aqueous (water) solution

• $Δ$ over the reaction arrow means heat is needed.

Properties and Changes

• Physical Properties:

  • Observed without changing a substance (e.g., color, odor, density, melting point).

• Chemical Properties:

  • Observed when a substance changes into another (e.g., flammability, oxidation).

Physical vs. Chemical Changes

• Physical Changes:

  • Easily reversible, no new products, often just a state change (e.g., ice melting).

• Chemical Changes:

  • Not easily reversed, new product(s) formed, reactants used up, often heat/light/sound/fizzing occurs (e.g., wood burning).

Simple Patterns of Chemical Reactivity

• Combination reactions

• Decomposition reactions

• Combustion reactions

• Precipitation reactions

• Oxidation-reduction reactions

Combination Reactions

• Two or more substances react to form one product.

Decomposition Reactions

• One substance breaks down into two or more substances.

• Example: Heating a metal carbonate produces carbon dioxide and a metal oxide: $CaCO<em>3(s) \rightarrow CaO(s) + CO</em>2(g)$

Combustion Reactions

• Rapid reactions that produce a flame, often involving oxygen.

• Burning compounds with C and H yields $CO<em>2$ and $H</em>2O$.

Precipitation Reactions

• Two solutions containing soluble salts mix to produce an insoluble salt (precipitate).

Oxidation-Reduction Reactions (Redox)

• Loss of electrons is oxidation.

• Gain of electrons is reduction.

• One cannot occur without the other.

Oxidation Numbers

• Used to determine if a redox reaction occurred.

Rules to Assign Oxidation Numbers

1. Atoms in elemental form have an oxidation number of zero (e.g., C, $O_2$, Mg).

2. Monatomic ion oxidation number equals its charge (e.g., $Ca^{2+}$, $O^{2-}$, $Na^+$, $Cl^-$.

3. Nonmetals usually have negative oxidation numbers, but can be positive:

   • Oxygen: −2 (except in peroxide ion $O_2^{2-}$, where it is −1).

   • Hydrogen: +1 when bonded to a nonmetal, −1 when bonded to a metal.

   • Halogens: −1, unless combined with oxygen (oxyanions), where they are positive.

4. Sum of oxidation numbers in a neutral compound is zero; in a polyatomic ion, it equals the ion's charge.

Formula Weight (FW)

• Sum of atomic weights in a chemical formula.

• For an element, it's the atomic weight from the periodic table.

• For an ionic compound, use the empirical formula.

• Example: $FW (H<em>2SO</em>4) = 2(1.0 \text{ amu}) + 1(32.1 \text{ amu}) + 4(16.0 \text{ amu}) = 98.1 \text{ amu}$

Molecular Weight (MW)

• Formula weight for a molecule.

• Sum of atomic weights of atoms in a molecule.

• Example: $MW (C<em>6H</em>{12}O_6) = 6(12.0 \text{ amu}) + 12(1.0 \text{ amu}) + 6(16.0 \text{ amu}) = 180.0 \text{ amu}$