Chapter 4: Chemical Reactions and Chemical Quantities – Study Notes

Chapter 4: Chemical Reactions and Chemical Quantities

Chemical Reactions

Chemical reactions are fundamental processes in chemistry where substances are transformed into new substances with different properties. Understanding chemical reactions is essential for predicting the outcomes of chemical changes and for quantifying the substances involved.

• Chemical reaction: A process in which one or more substances are converted into one or more different substances, involving chemical changes in matter and resulting in new chemical substances.

• Combustion reaction: A specific type of chemical reaction where a substance combines with oxygen to form one or more oxygen-containing compounds, typically emitting heat.

• Chemical equations: Represent the transformation of reactants into products, showing formulas, states, and relative numbers of molecules.

States of Reactants and Products

The physical state of reactants and products is indicated in chemical equations and can influence the type of reaction that occurs.

Abbreviation	State
(g)	Gas
(l)	Liquid
(s)	Solid
(aq)	Aqueous (water solution)

• Solid product: Indicates a precipitation reaction.

• Gas product: Indicates a gas evolution reaction (formation of bubbles).

• Aqueous reactions: Occur in water, very common in chemistry.

• Combustion reactions: Involve burning a compound in oxygen.

Types of Chemical Equations

Chemical reactions can be classified into several types based on the nature of the reactants and products.

• Combustion reactions: Organic compound + O2 → CO2 + H2O. Example: $C_3H_8 + 5O_2 \rightarrow 3CO_2 + 4H_2O$

• Combination reactions: Two or more reactants form one product. Example: $N_2 + 3H_2 \rightarrow 2NH_3$

• Decomposition reactions: One reactant forms multiple products. Example: $2KClO_3 \rightarrow 2KCl + 3O_2$

• Net ionic reactions, Acid-Base reactions: Covered in later chapters.

Organic compounds contain carbon and hydrogen, and possibly oxygen.

Combustion Reactions

Combustion reactions are important for energy production and involve burning hydrocarbons or other organics in air or pure oxygen.

• Unbalanced reaction: $CH_4 (g) + O_2 (g) \rightarrow CO_2 (g) + H_2O (g)$

• Balanced reaction: $CH_4 (g) + 2O_2 (g) \rightarrow CO_2 (g) + 2H_2O (g)$

• Balancing the equation satisfies the law of conservation of mass.

Balancing Chemical Equations

Balancing chemical equations ensures that the number and type of atoms are the same on both sides, reflecting the conservation of mass.

• Conservation of mass: Matter can neither be created nor destroyed.

• Stoichiometric coefficients: Integers before molecular formulas indicating the relative number of molecules or moles.

• A balanced equation predicts the main reaction, though side reactions may occur in practice.

Steps for Balancing Chemical Equations

• Write out the total number of atoms of each element on both sides of the equation.

• Balance one element at a time, starting with elements that appear in only one substance on each side.

• Balance the remaining elements one at a time.

• Multiply through by appropriate integers to remove any fractions.

Example: $C_3H_8 + O_2 \rightarrow CO_2 + H_2O$

Reaction Stoichiometry

Stoichiometry uses the coefficients in a balanced chemical equation to determine the relative amounts of reactants and products in moles.

• Example reaction: $2C_8H_{18}(l) + 25O_2(g) \rightarrow 16CO_2(g) + 18H_2O(g)$

• 2 molecules (or moles) of octane react with 25 molecules (or moles) of oxygen to form 16 molecules (or moles) of carbon dioxide and 18 molecules (or moles) of water.

Using Stoichiometry

Stoichiometry allows calculation of the amount of product formed from a given amount of reactant using the balanced equation.

• Stoichiometric ratio: The ratio of coefficients from the balanced equation.

• Example: From $2C_8H_{18} \rightarrow 16CO_2$, the ratio is $\frac{16\ mol\ CO_2}{2\ mol\ C_8H_{18}} = \frac{8\ mol\ CO_2}{1\ mol\ C_8H_{18}}$

Converting Between Moles

To convert between moles of different substances in a reaction, use the stoichiometric coefficients from the balanced equation.

• General formula:

$n_A A \rightarrow n_B B$

$\text{Moles B} = \frac{n_B}{n_A} \times \text{Moles A}$

• $n_B/n_A$ is the stoichiometric ratio.

Mole-to-Mole Calculations

These calculations determine the amount of one substance produced or consumed from a known amount of another.

• Example: If 22.0 moles of $C_8H_{18}$ are burned, how many moles of $CO_2$ are formed?

$22.0\ mol\ C_8H_{18} \times \frac{16\ mol\ CO_2}{2\ mol\ C_8H_{18}} = 176\ mol\ CO_2$

• Similarly, calculate moles of $H_2O$ and $O_2$ used:

$22.0\ mol\ C_8H_{18} \times \frac{18\ mol\ H_2O}{2\ mol\ C_8H_{18}} = 198\ mol\ H_2O$

$22.0\ mol\ C_8H_{18} \times \frac{25\ mol\ O_2}{2\ mol\ C_8H_{18}} = 275\ mol\ O_2$

Mass-to-Mole-to-Mole-to-Mass Calculations

To determine the mass of a product from a given mass of reactant, convert mass to moles, use stoichiometry, then convert back to mass.

• Example: What mass of $H_2O$ is produced from 1.00 g $C_8H_{18}$?

$2C_8H_{18}(l) + 25O_2(g) \rightarrow 16CO_2(g) + 18H_2O(g)$

Molar masses: $MM(H_2O) = 18.016\ g/mol$, $MM(C_8H_{18}) = 114.224\ g/mol$

$1.00\ g\ C_8H_{18} \times \frac{1\ mol\ C_8H_{18}}{114.224\ g\ C_8H_{18}} \times \frac{18\ mol\ H_2O}{2\ mol\ C_8H_{18}} \times \frac{18.016\ g\ H_2O}{1\ mol\ H_2O} = 1.42\ g\ H_2O$

Limiting Reactants

In reactions with more than one reactant, the limiting reactant is the one that is completely consumed first, thus determining the maximum amount of product formed.

• Other reactants are present in excess.

• The reaction stops when the limiting reactant is used up.

• Example (analogy): Making pizza with limited ingredients; the ingredient that runs out first limits the number of pizzas made.

Steps for Limiting Reactant Problems

• Balance the chemical equation.

• Calculate the molar mass of each reactant.

• Calculate the mass of product from each reactant (assuming excess of the other).

• Identify which reactant produces the smallest mass of product; this is the limiting reactant.

• The other reactant is present "in excess".

Yield of a Reaction

Theoretical yield is the maximum amount of product that can be formed from given reactants, assuming complete reaction. Actual yield is the amount actually obtained, which is often less due to side reactions or incomplete conversion.

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

• Example: If the theoretical yield of $CO_2$ is 12.74 g, but only 10.01 g is produced:

$\text{yield} = \frac{10.01\ g\ CO_2}{12.74\ g\ CO_2} \times 100\% = 78.57\%$

Additional info: These notes expand on the original slides by providing definitions, stepwise procedures, and academic context for each concept, making them suitable for exam preparation and self-study.