Liquids and Solids, Properties of Solutions, Kinetics, Equilibrium, Acids and Bases, Solubility, Thermodynamics, Electrochemistry, Transition Metals

Flashcards covering Vocabulary for all the topics in the notes

Phase Changes

Changes of state from solid to liquid to gas that do not involve change in molecules, but a change in forces among molecules; examples include melting, evaporation/vaporization/boiling, sublimation, condensation, freezing/fusion, and deposition.

Intramolecular Forces

Chemical bonds (coulombic forces) that hold atoms together to form molecules through sharing electrons (covalent bonding), ionic bonding, or metallic bonding.

Intermolecular Forces

Forces that hold molecules (not atoms) together to form solids and liquids; include Dipole-Dipole forces, Hydrogen Bonding, and London Dispersion forces.

Dipole-Dipole Forces

Attraction between opposite partially charged ends of polar molecules; only 1% as strong as covalent or ionic bonds.

Hydrogen Bonding

Attraction between partially positively charged H atom attached to a highly electronegative atom (N, O, F) and a lone pair on another atom; a very strong dipole-dipole force.

London Dispersion Forces

Attraction between an instantaneous dipole and a dipole induced in a neighboring atom or particle; occurs in all atoms and molecules.

Polarizability

How easy the electron cloud can be distorted to produce temporary dipolar charge distribution; increases with the number of electrons.

Surface Tension

Resistance of a liquid to an increase in its surface area; increases with stronger intermolecular forces.

Capillary Action

Spontaneous rising of a liquid in a narrow tube involving adhesive and cohesive forces.

Viscosity

A liquid's resistance to flow; increases with larger intermolecular forces.

Vapor Pressure

Pressure exerted by a vapor; increases with temperature.

ΔH(vapor)

Energy required to vaporize 1 mole of a liquid at a pressure of 1 atm; also known as heat of vaporization or enthalpy of vaporization.

Equilibrium Vapor Pressure

Pressure of vapor present at equilibrium in a closed system.

Clausius-Clapeyron Equation

Relates vapor pressure and temperature; used to determine P(vap) at another temperature when ΔH(vap) and P(vap) at one temperature are known.

Heat of Fusion (ΔH(fus))

Enthalpy (energy) change that occurs at freezing point; endothermic.

Normal Melting Point

Temperature at which solid and liquid are at equilibrium with total pressure = 1 atm.

Normal Boiling Point

Temperature at which the vapor pressure of the liquid is exactly 1 atm.

Phase Diagram

Indicates phase at each temperature/pressure combination; lines represent phase transitions.

Triple Point

Temperature and pressure at which all three phases exist in a closed system (equilibrium).

Critical Point

Temperature and pressure at which gas and liquid are no longer different.

Solution

A homogenous mixture of a solute dissolved in a solvent.

Molarity (M)

Moles of solute per liter of solution.

Mass Percent

Percent by mass of the solute in the solution.

Mole Fraction (x)

Moles of a substance per total moles of mixture.

Molality (m)

Number of moles of solute per kilogram of solvent; does not change with temperature.

Enthalpy (Heat) of Hydration

Enthalpy change associated with the dispersal of a solute in water.

Henry's Law

Amount of gas dissolved is directly proportional to pressure of gas above solution (C = kP).

Colligative Property

Property of solutions that depends only on the amount of solute present, not its identity.

Raoult's Law

The vapor pressure of a solution is directly proportional to the mole fraction of solvent present.

Ideal Solutions

Solutions that obey Raoult's Law; Solute-solute, solvent-solvent, and solute-solvent interactions are similar.

Boiling Point Elevation

The increase in boiling point of a solvent due to the presence of a nonvolatile solute.

Freezing Point Depression

The decrease in freezing point of a solvent due to the presence of a solute.

Osmosis

Flow of solvent into a solution through a semipermeable membrane.

Osmotic Pressure (Pie)

The excess pressure on the solution that arises because of the difference in the liquid levels; Pie=MRT.

Isotonic Solutions

Solutions with identical osmotic pressures.

Van't Hoff Factor (i)

Number of particles a solute breaks into; i = moles of particles in solution / moles of solute dissolved.

Chemical Kinetics

The study of how fast a reaction proceeds.

Reaction Rate

Change in a concentration of a reactant or product per unit time; Molarity per second.

Rate Law

Expresses the rate of a reaction in terms of the concentrations of reactants; Rate=K[A]^m[B]^n.

Rate Constant (k)

A constant of proportionality between the reaction rate and the concentrations of reactants that appear in the rate law.

Differential Rate Law

Expresses how the rate depends on concentration.

Integrated Rate Law

Expresses how concentration depends on time.

Half-Life

Amount of time required for a reactant to reach half its original amount.

Reaction Mechanism

Series of elementary steps by which a chemical reaction occurs.

Elementary Step

Reaction whose rate law can be written from its molecularity.

Molecularity

Number of species that collide to produce reaction indicated by the step (Unimolecular, Bimolecular, Termolecular).

Intermediate

Species formed in an early step and consumed in a later step.

Catalyst

Species consumed in an early step and reformed in a later step; speeds up reaction.

Rate-Determining Step

The slowest step in a multistep reaction; determines rate law of reaction.

Activation Energy (Ea)

Minimum energy required to initiate chemical reaction.

Heterogeneous Catalysis

Catalyst is in a different phase than reactants.

Homogeneous Catalysis

Catalyst is in the same phase as reactants.

Chemical Equilibrium

State in which there are no observable changes as time goes by.

Law of Mass Action

Writing equilibrium constant expression.

Kc

Equilibrium constant based on concentration (M).

Kp

Equilibrium constant based on pressure (atm).

Equilibrium Position

Set of equilibrium concentrations; has infinite number of equilibrium positions.

Reaction Quotient, Q

Law of mass action using any concentrations instead of equilibrium concentrations.

Le Chatelier's Principle

If an external stress is applied to a system at equilibrium, the system adjusts in a way to reduce the stress and reach a new equilibrium.

Arrhenius Acid

Produces H+ in H2O.

Arrhenius Base

Produces OH- in H2O.

Bronsted-Lowry Acid

Proton (H+) donor.

Bronsted-Lowry Base

Proton (H+) acceptor.

Lewis Acid

Electron pair acceptor.

Lewis Base

Electron pair donor.

Autoionization of Water

The equilibrium constant for the reaction where a molecule of a substance ionizes by reacting with another molecule of the same substance; denoted as Kw.

Acid Ionization Constant (Ka)

Represents the fraction of the original acid that has been ionized in solution; value reflects the strength of acid.

Base Ionization Constant (Kb)

Equilibrium constant that measures how much a base ionizes in water; value indicates the strength of the base.

pH

Measure of acidity; pH range: Neutral Ph= 7, Basic pH>7, Acidic Ph<7.

Percent Dissociation

Percent of a weak acid that has dissociated into ions.

Polyprotic Acids

Acids that can produce more than one proton; dissociate in a stepwise manner.

Salts

Ionic compounds; cation is conjugate acid of a base, anion is conjugate base of acid.

Common Ion Effect

Shift in equilibrium position that occurs because of addition of ion already involved in equilibrium reaction; presence of common ion suppresses ionization of a weak acid or base.

Buffered Solution

Resists changes in pH with addition of either acid or bases; contains a weak acid or weak base and the conjugate.

Titration

Determine amount of acid or base in a solution; solution of known concentration (titrant) delivered from a buret into unknown solution (analyte) until substance being analyzed is just consumed.

Equivalence Point

Stoiciometric point where amount of titrant has been added to exactly react with all analyte originally present, signaled by color change of an indicator.

Solubility

Ability of a substance to dissolve; = x.

First Law of Thermodynamics

Energy can be converted from one form to another but cannot be created nor destroyed.

Second Law of Thermodynamics

The entropy of the UNIVERSE increases in spontaneous process and remains unchanged in an equilibirum process.

Third Law of Thermodynamics

The entropy of a perfect crystalline substance is zero at the absolute zero of temperature.

Entropy (S)

Measure of the randomness or disorder of a system; ΔS = Sf - Si.

Spontaneous process

Occurs without outside intervention, thermodynamically favored.

Gibbs Free Energy (ΔG)

Measure of spontaneity taking into account both enthalpy and entropy, at constant temperature and pressure and is negative for spontaneous processes, positive for nonspontaneous processes and zero at equilibrium.

Oxidation

Loss of electrons (OIL).

Reduction

Gain of electrons (RIG).

Reducing Agent

Electron donor; reactant that loses electrons (aka oxidized).

Oxidizing Agent

Electron acceptor; reactant gains electrons (aka reduced).

Cell Potential

Energy difference between half-reactions; also called electromotive force (emf).

Coordination Number

Number of bonds formed between metal ion and ligands in a complex ion.

Ligand

Neutral molecule or ion with lone electron pair (Lewis base) that can be used to form a bond with a metal ion.

Isomers

structural and stereoisomers (Optical, Geometric, Coordination, Linkage)