Chemical Reactions and Reaction Stoichiometry

Chemical Equations and Balancing

• Chemical Equations: Represent reactions by separating reactants (starting materials) on the left from products (ending materials) on the right using arrows.

• Law of Conservation of Mass: Equations must be balanced by changing coefficients (the number in front of a formula), never subscripts (the numbers within a formula), as changing subscripts alters the identity of the substance.

• State Symbols: Phases of matter are indicated in parentheses:

  • $(g)$ = gas

  • $(l)$ = liquid

  • $(s)$ = solid

  • $(aq)$ = aqueous (dissolved in water)

Patterns of Chemical Reactivity

• Combination Reactions: Two or more substances combine to form one product ($A + B \rightarrow C$). Metals and nonmetals react to form ionic compounds based on common group charges.

• Decomposition Reactions: A single reactant breaks down into two or more substances ($C \rightarrow A + B$).

  • Metal Carbonates: Decompose when heated to produce a metal oxide and carbon dioxide ($CO_2$).

• Combustion Reactions: Rapid reactions usually involving oxygen ($O_2$) as a reactant. The combustion of hydrocarbons (compounds containing $C$ and $H$) always produces $CO_2$ and $H_2O$.

Formula Weight, Molecular Weight, and the Mole

• Formula Weight (FW): The sum of atomic weights of atoms in a chemical formula (e.g., $23.0\,amu$ for $Na$). Used for all compounds, including ionic ones using empirical formulas.

• Molecular Weight (MW): The formula weight for a molecule (e.g., $180.0\,amu$ for glucose, $C_6H_{12}O_6$).

• Avogadro’s Number: $6.02 \times 10^{23}$ particles (atoms, ions, or molecules) equals one mole ($mol$).

• Molar Mass: The mass in grams of one mole of a substance ($g/mol$). It is numerically equal to the formula weight in $amu$.

• Percent Composition: Calculated to find the mass percentage of each element in a compound:   \text{% Element} = \frac{(\text{number of atoms}) \times (\text{atomic weight})}{\text{formula weight}} \times 100\%.

Empirical and Molecular Formulas

• Determining Empirical Formulas:

  1. Assume a $100\,g$ sample (convert mass % to grams).

  2. Convert grams to moles using molar mass.

  3. Calculate the simplest whole-number mole ratio by dividing by the smallest number of moles.

• Molecular Formulas: Found by determining the whole-number multiple relating the empirical formula weight to the actual molar mass.

• Combustion Analysis: Used for compounds with $C$, $H$, and $O$. Mass of $C$ is derived from $CO_2$, mass of $H$ from $H_2O$, and mass of $O$ is the difference between the sample mass and the sum of $C$ and $H$.

Stoichiometry and Yield

• Stoichiometric Calculations: Coefficients in balanced equations provide the mole ratio needed to convert between amounts of different reactants and products.

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical Yield: The maximum amount of product that can be produced based on stoichiometry.

• Actual Yield: The amount of product actually obtained and measured in a lab.

• Percent Yield: The ratio of actual yield to theoretical yield:   $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

• Heat: Often denoted by the Greek symbol delta ($\Delta$) over the reaction arrow; heat amounts also depend on stoichiometry.