Physical Chemistry Lecture Notes

Flashcards for vocabulary review of key physical chemistry concepts.

Physical Chemistry

The study of the physical principles of chemistry.

Matter

Anything that occupies space and has mass.

Dimensional Analysis

A method of problem-solving using units to guide calculations.

SI Units

The International System of Units, a standard system of measurement.

Mole (mol)

The amount of a substance that contains as many elementary entities as there are atoms in exactly 12.00 grams of 12C.

Avogadro's Number (NA)

6.0221367 x 10^23, the number of elementary entities in one mole.

Molar Mass

The mass of one mole of a substance in grams.

Chemical Equation

A symbolic representation of a chemical reaction using chemical formulas.

Balancing Chemical Equations

Adjusting coefficients in a chemical equation to ensure the number of atoms of each element is the same on both sides.

Limiting Reagent

The reactant that is completely consumed in a chemical reaction.

Excess Reagent

The reactant that remains after the limiting reagent is completely consumed.

Theoretical Yield

The amount of product that would result if all the limiting reagent reacted.

Actual Yield

The amount of product actually obtained from a reaction.

Percent Yield

The ratio of actual yield to theoretical yield, multiplied by 100%.

Gases

Assume the volume and shape of their containers and are the most compressible state of matter.

Boyle's Law

P1 x V1 = P2 x V2, at constant temperature and amount of gas.

Charles' Law

V1/T1 = V2/T2, at constant pressure.

Avogadro's Law

V1 / n1 = V2 / n2, at constant temperature and pressure.

Ideal Gas Equation

PV = nRT, relates pressure, volume, moles, and temperature of an ideal gas.

Standard Temperature and Pressure (STP)

0°C (273.15 K) and 1 atm.

Dalton's Law of Partial Pressures

The total pressure of a mixture of gases is the sum of the partial pressures of the individual gases.

Kinetic Molecular Theory of Gases

A model describing the behavior of gases based on the motion and properties of gas molecules.

Gas Diffusion

The gradual mixing of molecules of one gas with molecules of another.

Gas Effusion

The process by which gas under pressure escapes from one compartment to another through a small opening.

Thermochemistry

The study of heat change in chemical reactions.

Thermodynamics

The scientific study of the interconversion of heat and other kinds of energy.

System

The specific part of the universe that is of interest in the study.

Surroundings

The rest of the universe outside the system.

Exothermic Process

Any process that gives off heat – transfers thermal energy from the system to the surroundings.

Endothermic Process

Any process in which heat has to be supplied to the system from the surroundings.

State functions

Properties that are determined by the state of the system, regardless of how that condition was achieved.

First Law of Thermodynamics

Energy can be converted from one form to another but cannot be created or destroyed.

Enthalpy (H)

Used to quantify the heat flow into or out of a system in a process that occurs at constant pressure.

Thermochemical Equations

Show the enthalpy changes as well as the mass relationships.

Specific Heat

The amount of heat (q) required to raise the temperature of one gram of the substance by one degree Celsius.

Heat Capacity (C)

The amount of heat (q) required to raise the temperature of a given quantity (m) of the substance by one degree Celsius.

Standard Enthalpy of Formation

Is the heat change that results when one mole of a compound is formed from its elements at a pressure of 1 atm.

Standard Enthalpy of Reaction

Is the enthalpy of a reaction carried out at 1 atm.

Hess's Law

When reactants are converted to products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.

Average Bond Energy

Is the enthalpy change required to break a particular bond in one mole of gaseous molecules.

Entropy (S)

Can be thought of as a measure of the randomness or disorder of a system.

Second Law of Thermodynamics

The entropy of the universe increases in a spontaneous process and remains unchanged in an equilibrium process.

Gibbs Free Energy (G)

Relates S, H and T of a system and can be used to predict spontaneity.

Standard Free-Energy of reaction (G0)

is the free-energy change for a reaction when it occurs under standard-state conditions.

Chemical equilibrium

Is achieved when the rates of the forward and reverse reactions are equal, and the concentrations of the reactants and products remain constant

Homogenous equilibrium

Applies to reactions in which all reacting species are in the same phase.

Heterogenous equilibrium

Applies to reactions in which reactants and products are in different phases.

Le Châtelier’s Principle

If an external stress is applied to a system at equilibrium, the system adjusts in such a way that the stress is partially offset as the system reaches a new equilibrium position.

Intermolecular Forces

Attractive forces between molecules.

Surface tension

The amount of energy required to stretch or increase the surface of a liquid by a unit area.

Cohesion

The intermolecular attraction between like molecules.

Adhesion

An attraction between unlike molecules

Viscosity

A measure of a fluid’s resistance to flow.

Electrochemistry

Is the study of the interconversion of electrical and chemical energy.

Oxidation

Loss of electrons.

Reduction

Gain of electrons.

Galvanic Cell

Salt Bridge-transports cations and anions to balance charge and completes the circuit; Anode Oxidation reaction e-; Cathode Reduction reaction.

Electrolysis

Is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur.