Chemical Reactions and Reaction Stoichiometry: Study Notes

Chemical Reactions and Reaction Stoichiometry

Chemical Equations

Chemical equations are the symbolic representation of chemical reactions, showing the relationships between reactants and products. They allow chemists to communicate the details of chemical changes in a concise format.

• Reactants are the starting substances, written on the left side of the equation.

• Products are the substances formed, written on the right side.

• An arrow (→) separates reactants from products.

• A plus sign (+) separates multiple reactants or products.

Example: $\mathrm{2H_2 + O_2 \rightarrow 2H_2O}$

Balancing Chemical Equations

Balancing chemical equations ensures the Law of Conservation of Mass is obeyed, meaning the number of atoms of each element is the same on both sides of the equation.

• Balance by changing coefficients, not subscripts.

• Start with elements that appear in only one reactant and one product.

• Continue balancing other elements, checking all at the end.

Example:

• Unbalanced: $\mathrm{CH_4 + O_2 \rightarrow CO_2 + H_2O}$

• Balanced: $\mathrm{CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O}$

Why Use Coefficients Instead of Changing Subscripts?

Changing subscripts alters the chemical identity of a compound. Coefficients adjust the number of molecules without changing their composition.

• Example: $\mathrm{2H_2(g) + O_2(g) \rightarrow 2H_2O(l)}$ (water)

• Example: $\mathrm{H_2(g) + O_2(g) \rightarrow H_2O_2(l)}$ (hydrogen peroxide)

Do not change the formula itself; only adjust the number of molecules.

Other Symbols in Chemical Equations

Additional symbols provide information about the physical states and reaction conditions.

• (g) = gas

• (l) = liquid

• (s) = solid

• (aq) = aqueous (dissolved in water)

• Symbols above the arrow (e.g., $\Delta$) indicate conditions such as heat.

Example: $\mathrm{CH_4(g) + 2O_2(g) \xrightarrow{\Delta} CO_2(g) + 2H_2O(g)}$

Simple Patterns of Chemical Reactivity

Chemical reactions can be classified into several types based on their patterns. The most common types are:

• Combination reactions

• Decomposition reactions

• Combustion reactions

Combination Reactions

In a combination reaction, two or more substances react to form a single product.

• General form: $\mathrm{A + B \rightarrow AB}$

• Common for elements forming compounds.

Combination Reaction	Description
$\mathrm{C + O_2 \rightarrow CO_2}$	Two or more reactants combine to form a single product.
$\mathrm{N_2 + 3H_2 \rightarrow 2NH_3}$	Elements react to form compounds.
$\mathrm{CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq)}$	Metal oxide reacts with water to form a hydroxide.

Combination Reaction Predictions: Metal and Nonmetal

When a metal reacts with a nonmetal, the product is typically an ionic compound. Predict the product using common charges for the elements involved.

• Example: $\mathrm{2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)}$

• Alkali metals react vigorously with halogens to form salts.

• Alkali metals also react with water: $\mathrm{2M(s) + 2H_2O(l) \rightarrow 2M^+(aq) + 2OH^-(aq) + H_2(g)}$

Combination Reaction Predictions: Nonmetals

• Halogens react with hydrogen to form hydrogen halides: $\mathrm{H_2(g) + X_2 \rightarrow 2HX(g)}$

• Interhalogen compounds form between different halogens: $\mathrm{Br_2(l) + F_2(g) \rightarrow 2BrF(g)}$

• Compounds between nonmetals typically contain covalent bonds.

Decomposition Reactions

In a decomposition reaction, a single compound breaks down into two or more simpler substances.

• General form: $\mathrm{AB \rightarrow A + B}$

• Often requires heat, light, or electricity.

Decomposition Reaction	Description
$\mathrm{2KClO_3(s) \rightarrow 2KCl(s) + 3O_2(g)}$	Potassium chlorate decomposes to potassium chloride and oxygen.
$\mathrm{CaCO_3(s) \rightarrow CaO(s) + CO_2(g)}$	Metal carbonates decompose to metal oxides and carbon dioxide.

Combustion Reactions

Combustion reactions are rapid reactions with oxygen that produce a flame. They are common for organic compounds.

• General form: Hydrocarbon + $\mathrm{O_2}$ → $\mathrm{CO_2}$ + $\mathrm{H_2O}$

• Products are usually carbon dioxide and water.

Example: $\mathrm{CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g)}$

Stoichiometry and Quantitative Information from Balanced Equations

Stoichiometry is the calculation of reactants and products in chemical reactions using balanced equations.

• Coefficients indicate relative numbers of molecules and moles.

• Mole ratios from coefficients are used to convert between substances.

Example: $\mathrm{4Fe(s) + 3O_2(g) \rightarrow 2Fe_2O_3(s)}$

To find moles of $\mathrm{Fe_2O_3}$ produced from 31.2 mol of Fe:

• Use the mole ratio: $\frac{2\text{ mol } Fe_2O_3}{4\text{ mol } Fe}$

Stoichiometric Calculations

Stoichiometric calculations involve converting between mass and moles using molar mass and mole ratios from the balanced equation.

• Convert grams to moles using molar mass.

• Use mole ratios to relate reactants and products.

• Convert moles back to grams if needed.

Example: How many grams of water can be produced from 1.00 g of $\mathrm{C_6H_{12}O_6}$?

• Balanced equation: $\mathrm{C_6H_{12}O_6(s) + 6O_2(g) \rightarrow 6CO_2(g) + 6H_2O(l)}$

• Convert mass of $\mathrm{C_6H_{12}O_6}$ to moles, use mole ratio, then convert moles of $\mathrm{H_2O}$ to grams.

Limiting Reactants

In reactions where reactants are not mixed in exact stoichiometric ratios, the limiting reactant is the one that is completely consumed first and determines the amount of product formed.

• Identify the limiting reactant by comparing mole ratios.

• Use the limiting reactant for all stoichiometric calculations.

Example: $\mathrm{2H_2(g) + O_2(g) \rightarrow 2H_2O(g)}$

If less $\mathrm{H_2}$ is present than required, it is the limiting reactant.

Theoretical Yield and Percent Yield

Theoretical yield is the maximum amount of product that can be formed from the limiting reactant, calculated using stoichiometry. Actual yield is the amount obtained in practice.

• Percent yield compares actual yield to theoretical yield:

$\text{Percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100$

Example: If 42.8 kg of Ti is produced from a reaction with a theoretical yield of 50 kg, percent yield is:

$\frac{42.8}{50} \times 100 = 85.6\%$

Additional info: These notes expand on the provided slides and images, filling in missing context and explanations for clarity and completeness.