Chemical Quantities and Chemical Reactions: Study Notes

Chemical Quantities and Reactions

The Mole Concept

The mole is a fundamental unit in chemistry used to express amounts of a chemical substance. It allows chemists to count atoms, molecules, ions, or formula units by weighing them.

• Definition: 1 mole (mol) contains $6.02 \times 10^{23}$ particles (Avogadro's number).

• Analogy: Just as a dozen means 12 items, a mole means $6.02 \times 10^{23}$ particles.

• Types of Particles: Atoms, molecules, ions, or formula units depending on the substance.

Example: 1 mole of water (H2O) contains $6.02 \times 10^{23}$ molecules of H2O.

Moles in Chemical Formulas

Chemical formulas indicate the number of moles of each element in one mole of a compound.

• Example: Sucrose (C12H22O11): - 1 mole of sucrose contains:

  • 12 moles of carbon (C)

  • 22 moles of hydrogen (H)

  • 11 moles of oxygen (O)

Molar Mass

Molar mass is the mass in grams of one mole of a substance. For elements, it is numerically equal to the atomic mass (in amu) but expressed in g/mol.

• Formula: $\text{Molar mass (g/mol)} = \text{Atomic mass (amu)}$

• Example: Carbon (C): 12.01 g/mol

Calculating Molar Mass of Compounds

1. Find the molar mass of each atom in the compound.

2. Multiply by the number of atoms of each element.

3. Add the total masses together.

Example: Calculate the molar mass of HClO3:

• H: 1.01 g/mol × 1 = 1.01 g

• Cl: 35.45 g/mol × 1 = 35.45 g

• O: 16.00 g/mol × 3 = 48.00 g

• Total: 1.01 + 35.45 + 48.00 = 84.46 g/mol

Conversions: Grams, Moles, and Particles

Conversions between mass, moles, and number of particles are fundamental in chemical calculations.

• Grams to Moles: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar mass (g/mol)}}$

• Moles to Grams: $\text{Mass (g)} = \text{Moles} \times \text{Molar mass (g/mol)}$

• Moles to Particles: $\text{Particles} = \text{Moles} \times 6.02 \times 10^{23}$

• Particles to Moles: $\text{Moles} = \frac{\text{Particles}}{6.02 \times 10^{23}}$

Example: How many molecules are in 2 moles of CO2? $2 \text{ mol} \times 6.02 \times 10^{23} = 1.204 \times 10^{24}$ molecules

Chemical Reactions and Equations

Types of Chemical Reactions

Chemical reactions can be classified into several types based on the rearrangement of atoms and molecules.

• Combination (Synthesis): $A + B \rightarrow AB$

• Decomposition: $AB \rightarrow A + B$

• Single Replacement: $AB + C \rightarrow AC + B$

• Double Replacement: $AB + CD \rightarrow AD + CB$

• Combustion: $C_xH_y + O_2 \rightarrow CO_2 + H_2O$

Example: Combustion of methane: $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$

Balancing Chemical Equations

Balancing 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.

• Adjust coefficients to balance atoms for each element.

• Never change subscripts in chemical formulas.

Oxidation and Reduction (Redox) Reactions

Redox reactions involve the transfer of electrons between substances.

• Oxidation: Loss of electrons (increase in oxidation state)

• Reduction: Gain of electrons (decrease in oxidation state)

Example: $2Al^{3+} + 3Sn \rightarrow 2Al + 3Sn^{2+}$ - Aluminum ions are reduced (gain electrons). - Tin is oxidized (loses electrons).

Stoichiometry and Mass Calculations

Stoichiometric Calculations

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

• Steps:

  1. Convert grams of a substance to moles using molar mass.

  2. Use the mole ratio from the balanced equation to find moles of another substance.

  3. Convert moles back to grams if required.

Example: If 10 g of A reacts with B, how many grams of product C are formed?

Energy in Chemical Reactions

Exothermic and Endothermic Reactions

Chemical reactions involve energy changes, usually in the form of heat.

• Exothermic Reaction: Releases heat; products have lower energy than reactants. $E_{\text{product}} < E_{\text{reactant}}$

• Endothermic Reaction: Absorbs heat; products have higher energy than reactants. $E_{\text{product}} > E_{\text{reactant}}$

• Energy Change: $\Delta H = E_{\text{product}} - E_{\text{reactant}}$

Example: Combustion of gasoline is exothermic; photosynthesis is endothermic.

Energy Diagrams

Energy diagrams show the progress of a reaction and the relative energies of reactants and products.

• Exothermic: Downward slope (energy released)

• Endothermic: Upward slope (energy absorbed)