States of Matter, Solution Chemistry, and Acids/Bases: CHEM 1150 Ch. 7, 8, & 9.1-9.5 Study Notes

States of Matter and Their Attractive Forces

Pressure and Gas Laws

Pressure is a fundamental concept in chemistry, describing the force exerted by particles against the walls of their container. Understanding pressure and its units is essential for studying gases and their behavior.

• Pressure: The force exerted against a given area.

• Atmosphere (atm): Common unit for measuring pressure, especially in chemistry.

• Pascal (Pa): SI unit of pressure. $1~\text{Pa} = 1~\text{N}/\text{m}^2$

• Pounds per square inch (psi): Measures pressure as the force (in pounds) applied to an area of 1 square inch. Atmospheric pressure at sea level is about 14.7 psi.

• Millimeters of mercury (mmHg): Derived from the height of mercury in a barometer. $1~\text{atm} = 760~\text{mmHg}$

Example: Atmospheric pressure at sea level is 1 atm, which is equivalent to 760 mmHg, 101.325 kPa, or 14.7 psi.

Molecular Theory of Gases

The kinetic molecular theory explains the behavior of gases based on the motion and energy of their particles.

• Gas particles are far apart; most of the volume is empty space.

• Particles are in constant, random motion and move rapidly.

• No significant attractive forces between gas particles.

• Kinetic energy is directly proportional to absolute temperature ($E_{\text{kinetic}} \propto T$).

Example: Increasing the temperature of a gas increases the speed of its particles.

Combined Gas Law

The combined gas law relates pressure, volume, and temperature for a fixed amount of gas.

• $\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$

• Allows calculation of changes when more than one variable changes.

Intermolecular Forces

Intermolecular forces are the attractive forces between molecules, influencing physical properties such as boiling point and solubility.

• London Dispersion Forces: Temporary attractive forces due to momentary uneven electron distribution. Present in all molecules, especially nonpolar ones.

• Dipole-Dipole Forces: Attraction between molecules with permanent dipoles (polar molecules).

• Hydrogen Bonding: Strong dipole-dipole interaction involving hydrogen bonded to O, N, or F. Requires a donor hydrogen and an acceptor pair of electrons.

Example: Water molecules exhibit hydrogen bonding, leading to high boiling point.

Boiling Point and Vapor Pressure

The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure.

• Vapor pressure: Pressure exerted by vapor above a liquid at equilibrium.

• Boiling point increases with stronger intermolecular forces.

Solution Chemistry—Sugar and Water Do Mix

Solutions and Solubility

A solution is a homogeneous mixture of two or more substances. Solubility describes how much solute can dissolve in a solvent at a given temperature.

• Solute: Substance present in smaller amount.

• Solvent: Substance present in larger amount.

• "Like dissolves like": Polar solvents dissolve polar solutes; nonpolar solvents dissolve nonpolar solutes.

Example: Sugar (polar) dissolves in water (polar), but oil (nonpolar) does not.

Types of Mixtures

• Solution: Homogeneous, transparent, particles < 1 nm.

• Colloid: Heterogeneous, particles between 1 and 100 nm, do not settle out.

• Suspension: Heterogeneous, particles > 1000 nm, can be separated by filtration or centrifugation.

Saturated and Unsaturated Solutions

• Unsaturated solution: Contains less solute than the maximum possible.

• Saturated solution: Contains the maximum amount of solute that can dissolve.

Solubility and Temperature

• Solubility of most solids increases with temperature.

• Solubility of gases decreases with increasing temperature.

Electrolytes and Nonelectrolytes

• Electrolytes: Compounds that produce ions in solution and conduct electricity.

• Strong electrolytes: Completely ionize in water (e.g., NaCl).

• Weak electrolytes: Partially ionize in water (e.g., acetic acid).

• Nonelectrolytes: Do not produce ions (e.g., sugar).

Concentration Units

• Percent concentration: Mass/mass (% m/m), volume/volume (% v/v), mass/volume (% m/v).

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

• Equivalents (Eq): Amount of ion that supplies 1 mole of charge.

Osmosis and Osmotic Pressure

Osmosis is the movement of water across a semipermeable membrane from low to high solute concentration. Osmotic pressure is the pressure required to stop this flow.

• Isotonic solutions have the same osmolarity as plasma (0.3 osmoles/L).

• IV solutions are designed to be isotonic to minimize osmosis in patients.

Diffusion and Transport

• Diffusion: Movement of molecules from high to low concentration.

• Passive diffusion: No energy required.

• Facilitated transport: Uses proteins to help molecules cross membranes.

• Active transport: Requires energy to move molecules against concentration gradient.

Acids, Bases, and Buffers in the Body

Arrhenius and Brønsted-Lowry Definitions

Acids and bases are defined by their behavior in water and their ability to donate or accept protons.

• Arrhenius Acid: Produces H+ ions in water.

• Arrhenius Base: Produces OH- ions in water.

• Brønsted-Lowry Acid: Proton donor.

• Brønsted-Lowry Base: Proton acceptor.

• Amphiprotic: Can act as both acid and base (e.g., water).

Strong and Weak Acids/Bases

• Strong acids: Completely ionize in water (e.g., HCl, HBr, HI, HNO3).

• Strong bases: Group 1 and heavy group 2 hydroxides (e.g., NaOH, KOH, Ba(OH)2).

• Weak acids/bases: Partially ionize (e.g., CH3COOH, NH3).

Naming Acids

• Binary acids: "hydro" + anion name + "ic acid" (e.g., HCl = hydrochloric acid).

• Oxoacids: Anion name ending in "-ate" becomes "-ic acid"; "-ite" becomes "-ous acid" (e.g., H2SO4 = sulfuric acid).

Neutralization Reactions

When a strong acid reacts with a strong base, the products are water and a salt.

• $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$

Antacids

• Used to neutralize excess stomach acid.

• Common antacids: aluminum hydroxide, magnesium hydroxide, carbonates.

• Reaction produces salt, water, and carbon dioxide.

Chemical Equilibrium

Reversible reactions reach equilibrium when the rates of forward and reverse reactions are equal.

• Equilibrium constant ($K$): $K = \frac{[\text{products}]}{[\text{reactants}]}$

• Only concentrations of gases and aqueous solutions are included.

Le Châtelier’s Principle

If a system at equilibrium is disturbed, it will shift to counteract the disturbance and restore equilibrium.

• Adding reactants shifts equilibrium to products.

• Adding products shifts equilibrium to reactants.

Conjugate Acid-Base Pairs

• Acids and bases exist in pairs that differ by one proton.

• Example: $\text{CH}_3\text{COOH} + \text{H}_2\text{O} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}_3\text{O}^+$

Ionization of Water and the pH Scale

• Water ionizes to produce equal amounts of H+ and OH-.

• $[\text{H}^+][\text{OH}^-] = 1 \times 10^{-14}$ at 25°C

• pH: $\text{pH} = -\log[\text{H}^+]$

• pH < 7: acidic; pH = 7: neutral; pH > 7: basic

Example: Blood pH is tightly regulated between 7.35 and 7.45.

Unit	Symbol	Equivalent Value
Atmosphere	atm	1 atm
Millimeters of mercury	mmHg	760 mmHg
Pounds per square inch	psi	14.7 psi
Pascals	Pa	101,325 Pa

*Additional info: Some explanations and examples have been expanded for clarity and completeness, including definitions and context for intermolecular forces, solution types, and acid/base behavior.*