Chemical Reactions and Chemical Quantities: Study Notes

CH. 4: Chemical Reactions and Chemical Quantities

What is a Chemical Reaction

A chemical reaction involves the transformation of one or more substances into different substances. Chemical reactions are fundamental to chemistry, as they describe how matter changes and interacts.

• Definition: A process in which chemical change occurs, resulting in the formation of new substances with different properties.

• Conservation of Mass: The total mass of reactants equals the total mass of products.

• Signs of a Chemical Reaction: Heat, light, sound, gas formation, color change, or precipitate formation.

• Chemical Equation: A shorthand way to represent a chemical reaction, showing reactants and products. Example: Reactants → Products

Example: Combustion of methane:

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

State of Reactants and Products	Abbreviation
Solid	(s)
Liquid	(l)
Gas	(g)
Aqueous	(aq)

Writing Balanced Equations

Balanced chemical equations ensure the conservation of mass and atoms. Each side of the equation must have the same number of each type of atom.

• Steps to Balance:

  1. Write the unbalanced equation.

  2. Count atoms of each element on both sides.

  3. Add coefficients to balance atoms.

• Practice Examples:

  • Decomposition of potassium chlorate: $\mathrm{2\ KClO_3(s) \rightarrow 2\ KCl(s) + 3\ O_2(g)}$

  • Reaction of iron and copper(II) sulfate: $\mathrm{2\ Fe(s) + 3\ CuSO_4(aq) \rightarrow Fe_2(SO_4)_3(aq) + 3\ Cu(s)}$

Reaction Stoichiometry

Stoichiometry is the study of quantitative relationships between reactants and products in a chemical reaction. It allows chemists to predict the amounts of substances consumed and produced.

• Coefficients: Indicate the relative number of moles of each substance.

• Mole-to-Mole Ratios: Used to convert between amounts of reactants and products.

  • Example: Combustion of propane $\mathrm{C_3H_8 + 5\ O_2 \rightarrow 3\ CO_2 + 4\ H_2O}$ 1 mole propane reacts with 5 moles oxygen to produce 3 moles carbon dioxide and 4 moles water.

• Avogadro's Number: $6.022 \times 10^{23}$ particles per mole.

Reaction Stoichiometry from Grams

Stoichiometric calculations often require converting mass to moles using molar mass, then using mole ratios from the balanced equation.

• Mass-to-Mass Conversion Steps:

  1. Convert mass of reactant to moles.

  2. Use mole ratio to find moles of desired product.

  3. Convert moles of product to mass.

• Example:

  • If 50.0 g of propane ($\mathrm{C_3H_8}$) is burned, how many grams of oxygen are needed? $\mathrm{C_3H_8 + 5\ O_2 \rightarrow 3\ CO_2 + 4\ H_2O}$

Limiting Reactant and Theoretical Yield

In reactions with more than one reactant, the limiting reactant is the one that is completely consumed first, thus limiting the amount of product formed. The theoretical yield is the maximum amount of product possible, based on the limiting reactant.

• Limiting Reactant: Consumed completely in the reaction.

• Excess Reactant: Present in greater quantity than needed.

• Theoretical Yield: Calculated from the amount of limiting reactant.

• Example:

  • $\mathrm{2\ Mg(s) + O_2(g) \rightarrow 2\ MgO(s)}$

  • If 42.5 g Mg reacts with 33.8 g O2, calculate the limiting reactant and theoretical yield of MgO.

Percent Yield

Percent yield measures the efficiency of a reaction by comparing the actual yield to the theoretical yield.

• Theoretical Yield: Maximum possible product.

• Actual Yield: Amount of product actually obtained.

• Percent Yield Formula: $\mathrm{Percent\ Yield = \frac{Actual\ Yield}{Theoretical\ Yield} \times 100}$

• Example:

  • If 70.5 g MgO is the theoretical yield and 56.8 g MgO is isolated, percent yield is: $\mathrm{Percent\ Yield = \frac{56.8}{70.5} \times 100 = 80.6\%}$

Combustion Reactions

Combustion reactions involve a substance reacting with oxygen to form one or more oxygen-containing compounds, often releasing energy as heat and light.

• General Form: Hydrocarbon + O2 → CO2 + H2O

• Examples:

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

  • Ethanol combustion: $\mathrm{C_2H_5OH(l) + 3\ O_2(g) \rightarrow 2\ CO_2(g) + 3\ H_2O(g)}$

  • Magnesium combustion: $\mathrm{2\ Mg(s) + O_2(g) \rightarrow 2\ MgO(s)}$

Reactions of Alkali Metals

Alkali metals (Group 1A) are highly reactive, especially with water and oxygen. Their reactivity increases down the group.

• Reaction with Oxygen (Combustion):

  • $\mathrm{4\ Na(s) + O_2(g) \rightarrow 2\ Na_2O(s)}$

• Reaction with Water:

  • $\mathrm{2\ Na(s) + 2\ H_2O(l) \rightarrow 2\ NaOH(aq) + H_2(g)}$

• Reaction with Halogens:

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

• Reactivity increases with each successive alkali metal (e.g., Li < Na < K).

Reactions of Halogens

Halogens (Group 7A) are reactive nonmetals that readily form compounds with metals and other nonmetals.

• Example:

  • $\mathrm{Cl_2(g) + 2\ NaBr(aq) \rightarrow 2\ NaCl(aq) + Br_2(l)}$

• Halogens undergo displacement reactions and form salts with alkali metals.

Summary Table: Types of Chemical Reactions

Type	General Form	Example
Combustion	Hydrocarbon + O2 → CO2 + H2O	$\mathrm{CH_4 + 2\ O_2 \rightarrow CO_2 + 2\ H_2O}$
Combination	A + B → AB	$\mathrm{2\ Mg + O_2 \rightarrow 2\ MgO}$
Decomposition	AB → A + B	$\mathrm{2\ KClO_3 \rightarrow 2\ KCl + 3\ O_2}$
Single Displacement	A + BC → AC + B	$\mathrm{Fe + CuSO_4 \rightarrow FeSO_4 + Cu}$
Double Displacement	AB + CD → AD + CB	$\mathrm{Na_2SO_4 + BaCl_2 \rightarrow BaSO_4 + 2\ NaCl}$

Additional info: These notes expand on the original slides by providing definitions, stepwise explanations, and extra examples for clarity and completeness.