General Chemistry Study Notes: Chemical Reactions, Intermolecular Forces, and Properties of Solutions

Chapter 4: Types of Chemical Reactions & Solution Stoichiometry

Chemical Reactions and Balancing Equations

Chemical reactions occur when substances, called reactants, are transformed into new substances, known as products. Balancing chemical equations ensures the conservation of mass by making sure the number of atoms of each element is the same on both sides of the equation.

• Balancing Equations: Adjust coefficients (whole numbers in front of formulas) to ensure equal numbers of each atom on both sides.

• Example: $2H_2 + O_2 \rightarrow 2H_2O$

Types of Chemical Reactions

• Precipitation Reaction: Two solutions mix and a solid (precipitate) forms. Example: $AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)$

• Acid-Base Reaction: Acid + Base → Water + Salt. Example: $HCl + NaOH \rightarrow NaCl + H_2O$

• Redox Reaction: Electrons are transferred between substances. Example: $Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu$

Stoichiometry

Stoichiometry uses balanced chemical equations to determine the amounts of reactants and products involved in a reaction. It involves converting between grams, moles, and molecules using molar ratios.

• Key Steps:

  1. Convert quantities to moles.

  2. Use mole ratios from the balanced equation.

  3. Convert moles to desired units (grams, molecules).

• Example: How many grams of $H_2O$ are produced from 4 grams of $H_2$?

Limiting Reactant

The limiting reactant is the substance that runs out first in a chemical reaction, thus limiting the amount of product formed.

• Example: If you have 5 mol $H_2$ and 2 mol $O_2$, $O_2$ is the limiting reactant in the formation of water.

Chapter 11: Intermolecular Forces, Liquids, and Solids

Intermolecular Forces (IMFs)

Intermolecular forces are the forces of attraction between molecules, which influence the physical properties of substances such as boiling and melting points.

• London Dispersion Forces: Weak forces present in all molecules, including nonpolar ones.

• Dipole-Dipole Forces: Occur in polar molecules due to the attraction between positive and negative ends.

• Hydrogen Bonding: A very strong IMF occurring when hydrogen is bonded to nitrogen, oxygen, or fluorine.

Properties of Liquids and Solids

• Boiling/Melting Points: Stronger IMFs result in higher boiling and melting points.

• Viscosity: The "thickness" or resistance to flow of a liquid. Higher IMFs lead to higher viscosity.

• Solids: Can be crystalline (ordered structure) or amorphous (random structure).

• Phase Changes: Melting, freezing, boiling, condensation, etc. These involve energy being absorbed or released.

Chapter 13: Properties of Solutions

Solutions and Concentration

A solution is a homogeneous mixture of a solute (the substance dissolved) in a solvent (the substance doing the dissolving).

• Concentration: Measures how much solute is present in a given amount of solvent.

  • Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

  • Other units: Molality, percent by mass, parts per million (ppm).

Solubility

Solubility refers to the maximum amount of solute that can dissolve in a solvent at a given temperature and pressure. It depends on the nature of both solute and solvent.

• Example: Sugar dissolves more in hot water than cold water.

Colligative Properties

Colligative properties depend on the number of particles dissolved in a solution, not their identity.

• Boiling Point Elevation: Adding solute increases the boiling point of the solvent. $\Delta T_b = iK_b m$

• Freezing Point Depression: Adding solute lowers the freezing point of the solvent. $\Delta T_f = iK_f m$

• Osmotic Pressure: Pressure required to prevent osmosis. $\Pi = iMRT$

Electrolytes

Electrolytes are substances that conduct electricity when dissolved in water, such as salts.

• Example: $NaCl$ in water dissociates into $Na^+$ and $Cl^-$ ions, allowing the solution to conduct electricity.