Chemical Reactions and Quantities: Study Notes for Introductory Chemistry

Chemical Reactions and Quantities

Learning Objectives

This section outlines the key skills and concepts students should master regarding chemical reactions and quantitative relationships in chemistry.

• Write balanced chemical equations from reactant and product formulas.

• Classify chemical reactions as combination, decomposition, single replacement, double replacement, or oxidation-reduction.

• Identify oxidation and reduction in reactions; determine which reactants are oxidized or reduced.

• Use Avogadro's number to calculate the number of particles in a given number of moles.

• Calculate molar mass from a chemical formula.

• Convert between grams and moles of a substance using molar mass.

• Apply stoichiometric relationships from balanced equations to calculate quantities of reactants and products.

• Identify the limiting reagent and calculate the amount of product formed and excess reactant remaining.

• Calculate heat changes for exothermic and endothermic reactions.

Chemical Changes and Evidence

Recognizing Chemical Changes

A chemical change occurs when a substance is transformed into one or more new substances with different properties and compositions.

• Evidence of chemical change includes:

  • Formation of bubbles (gas production)

  • Change in color

  • Production of a solid (precipitate)

  • Heat or light produced or absorbed

• Examples: Burning wood, rotting fruit, baking a cake, rusting metal, digestion, souring milk, exploding fireworks, cooking an egg.

Chemical Equations

Structure of a Chemical Equation

A chemical equation is a written representation of a chemical reaction, showing the reactants and products, their physical states, and the proportions in which they react.

• General format: $aA + bB \rightarrow cC + dD$

• Components:

  • Chemical formula: Indicates the identity of each substance.

  • Coefficients: Numbers in front of formulas showing the relative number of molecules or moles.

  • Subscripts: Numbers within formulas indicating the number of atoms in each molecule.

  • State symbols: (s) for solid, (l) for liquid, (g) for gas, (aq) for aqueous solution.

  • Reaction arrow: Separates reactants (left) from products (right).

• Example: $2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$

Types of Chemical Reactions

Classification of Reactions

Chemical reactions are classified based on the changes that occur to the reactants and products.

• Combination (Synthesis) Reaction: Two or more substances combine to form one product. Example: $2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)$

• Decomposition Reaction: One substance splits into two or more simpler substances. Example: $2HgO(s) \rightarrow 2Hg(l) + O_2(g)$

• Single Replacement Reaction: One element replaces another in a compound. Example: $Zn(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2(g)$

• Double Replacement Reaction: Two compounds exchange ions to form two new compounds. Example: $AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)$

• Combustion Reaction: A substance reacts with oxygen to produce energy (heat and light), often forming $CO_2$ and $H_2O$. Example: $CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g) + \text{energy}$

• Oxidation-Reduction (Redox) Reaction: Electrons are transferred between reactants; includes oxidation (loss of electrons) and reduction (gain of electrons).

Balancing Chemical Equations

Principles and Tips

Balancing chemical equations ensures the law of conservation of mass is obeyed: the number of atoms of each element is the same on both sides of the equation.

• Steps to balance:

  1. Write correct formulas for all reactants and products.

  2. Count the number of atoms of each element on both sides.

  3. Add coefficients to balance one element at a time.

  4. Check that all elements are balanced and coefficients are in the lowest ratio.

• Example: Balancing the synthesis of aluminum sulfide: Unbalanced: $Al(s) + S(s) \rightarrow Al_2S_3(s)$ Balanced: $2Al(s) + 3S(s) \rightarrow Al_2S_3(s)$

Oxidation-Reduction (Redox) Reactions

Electron Transfer and Identification

Redox reactions involve the transfer of electrons between substances. Oxidation is the loss of electrons, while reduction is the gain of electrons.

• Oxidation: Loss of electrons; increase in oxidation state.

• Reduction: Gain of electrons; decrease in oxidation state.

• Mnemonic: "OIL RIG" — Oxidation Is Loss, Reduction Is Gain (of electrons).

• Example: $2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)$ Sodium is oxidized (loses electrons), chlorine is reduced (gains electrons).

• Biological Example: FAD (flavin adenine dinucleotide) is reduced to FADH2 by gaining two hydrogen atoms and two electrons.

The Mole and Avogadro's Number

Counting Particles in Chemistry

The mole is the SI unit for measuring the amount of substance. One mole contains Avogadro's number of particles.

• Avogadro's Number: $6.02 \times 10^{23}$ particles per mole.

• Types of particles: Atoms, molecules, formula units, ions.

• Conversion factors:

  • $1 \text{ mole} = 6.02 \times 10^{23}$ particles

  • $\frac{1 \text{ mole}}{6.02 \times 10^{23} \text{ particles}}$ or $\frac{6.02 \times 10^{23} \text{ particles}}{1 \text{ mole}}$

• Example: Number of $CO_2$ molecules in $0.500$ mole: $0.500 \text{ mole} \times \frac{6.02 \times 10^{23} \text{ molecules}}{1 \text{ mole}} = 3.01 \times 10^{23}$ molecules

Molar Mass and Chemical Formulas

Calculating Molar Mass

Molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol). It is calculated by summing the atomic masses of all atoms in the chemical formula.

• Example: Molar mass of lithium carbonate ($Li_2CO_3$): $2 \times 6.94 \text{ g (Li)} + 12.01 \text{ g (C)} + 3 \times 16.00 \text{ g (O)} = 73.89 \text{ g/mol}$

• Conversion factors: $\frac{1 \text{ mole}}{\text{molar mass in g}}$ or $\frac{\text{molar mass in g}}{1 \text{ mole}}$

Stoichiometry: Quantitative Relationships in Reactions

Using Balanced Equations for Calculations

Stoichiometry involves using balanced chemical equations to calculate the amounts of reactants and products.

• Mole ratios: Derived from coefficients in the balanced equation.

• Example: $2Ag(s) + S(s) \rightarrow Ag_2S(s)$ Mole ratio: $2$moles Ag : $1$ mole S : $1$ mole Ag2S

• Mass relationships: Use molar mass to convert between grams and moles.

• Example: $737$g NaCl$\rightarrow$moles: $737 \text{ g} \times \frac{1 \text{ mole}}{58.44 \text{ g}} = 12.6 \text{ moles}$

Limiting Reactant and Percent Yield

Determining Maximum Product and Efficiency

The limiting reactant is the substance that is completely consumed first, limiting the amount of product formed. Percent yield measures the efficiency of a reaction.

• Limiting reactant: Calculate the amount of product from each reactant; the smaller amount determines the maximum product.

• Percent yield: $\text{Percent yield} = \frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100\%$

• Example: If $3.00$ moles CO and $5.00$ moles H2 react to form CH3OH: $CO(g) + 2H_2(g) \rightarrow CH_3OH(g)$ CO can produce $3.00$ moles CH3OH; H2 can produce $2.50$ moles CH3OH. H2 is limiting; maximum product is $2.50$ moles.

Conservation of Mass

Law of Conservation of Mass

In chemical reactions, mass is conserved: the total mass of reactants equals the total mass of products.

• Example: $2Ag(s) + S(s) \rightarrow Ag_2S(s)$ $2 \times 107.9 \text{ g (Ag)} + 32.07 \text{ g (S)} = 247.99 \text{ g (Ag}_2\text{S)}$

Summary Table: Types of Chemical Reactions

Type	Description	General Equation	Example
Combination	Two or more reactants form one product	$A + B \rightarrow AB$	$2Na + Cl_2 \rightarrow 2NaCl$
Decomposition	One reactant splits into two or more products	$AB \rightarrow A + B$	$2HgO \rightarrow 2Hg + O_2$
Single Replacement	One element replaces another in a compound	$A + BC \rightarrow AC + B$	$Zn + 2HCl \rightarrow ZnCl_2 + H_2$
Double Replacement	Exchange of ions between two compounds	$AB + CD \rightarrow AD + CB$	$AgNO_3 + NaCl \rightarrow AgCl + NaNO_3$
Combustion	Reaction with oxygen producing energy	$C_xH_y + O_2 \rightarrow CO_2 + H_2O$	$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$
Redox	Electron transfer between reactants	Varies	$2Na + Cl_2 \rightarrow 2NaCl$

Additional info: Some explanations and examples have been expanded for clarity and completeness, including biological redox reactions and percent yield calculations.