CHEM Unit 2

Lewis Theory

Works reasonably well for light atoms in main-group compounds, emphasizing the electron pair bond and octets.

Schrödinger Equation

A mathematical equation that describes the energy of a given state with potential and kinetic energy terms; useful for tackling chemical bonding.

Wavefunction

A solution to the Schrödinger equation that describes the behavior of an electron; squaring the wavefunction gives a probability distribution.

Atomic Orbitals

Characterized by quantum numbers including principal (n), angular momentum (l), magnetic (ml), and spin (ms). Phase is important.

Pauli Exclusion Principle

No two electrons can have the same four quantum numbers.

Molecular Orbital Theory

A delocalized view of chemical bonding.

Valence Bond Theory

A localized view of chemical bonding.

Molecular Orbitals

Solutions to the molecular problem, formed by combining atomic orbitals.

Symmetry

Molecular orbitals must have the same symmetry to combine.

Sigma (σ) symmetry

Symmetric with respect to rotation about the bond axis.

Pi (π) symmetry

Antisymmetric with respect to rotation about the bond axis.

Valence Electrons

The highest energy electrons in an atom, chemically accessible and focus of Lewis Theory.

Core Orbitals

Inner orbitals (e.g., 1s) that remain largely atom-centered and act as a screen for valence electrons; too low in energy to interact effectively with valence orbitals.

Bond Order

Indicates bond strength; calculated as (# bonding electrons - # antibonding electrons) / 2.

Paramagnetic

Substances with unpaired electrons that are attracted to a magnetic field.

Diamagnetic

Substances with all electrons paired, showing no attraction to a magnetic field.

Zeff (Effective Nuclear Charge)

The net positive charge experienced by an electron in a multi-electron atom; calculated as Z (number of protons) - S (shielding constant).

Orbital Mixing

The interaction between s and p orbitals, which can change the order of energy levels in molecular orbitals (e.g., in N2).

Heterodiatomics

Molecules composed of two different atoms, requiring consideration of the relative energies of the atomic orbitals.

Delocalization

The ability of electrons to move throughout a molecule rather than being localized between a pair of atoms.

Valence Bond Theory

Presents a localized view of bonding and focuses on electron pair bonds.

Valence Bond Orbital

Formed when two atomic orbitals come together.

sp3 Hybridized Orbitals

Four equivalent valence bond orbitals oriented at 109.5 degrees.

Sigma Bond

A single bond formed by the direct overlap of orbitals.

Pi Bond

A bond formed by the sideways overlap of p orbitals.

sp2 Hybridization

Hybridization involving one s and two p orbitals, resulting in three sp2 orbitals and one unhybridized p orbital.

sp Hybridization

Hybridization involving one s and one p orbital, resulting in two sp orbitals and two unhybridized p orbitals.

Suboctet

An exception to the octet rule where a central atom has fewer than eight electrons.

Octet Expansion

An exception to the octet rule where a central atom has more than eight electrons.

sp3d Hybridization

Hybridization that includes d-character and often results in an octahedral geometry.

Dipole Moment

Positive and negative charges separated by a distance, oriented towards the more electronegative atoms.

Molecular Dipole

The sum of bond dipoles, resulting in a molecular dipole moment.

Polar Molecules

Molecules with a net dipole moment due to uneven electron distribution.

Nonpolar Molecules

Molecules where bond dipoles cancel out, resulting in no net dipole moment.

Intermolecular Forces (IMFs)

Attractive or repulsive forces between molecules that influence physical properties.

Dispersive Forces (London Dispersion Forces)

Weak, temporary attractive forces between all molecules, caused by temporary fluctuations in electron distribution.

Dipole-Dipole Interactions

Attractive forces between polar molecules due to the interaction of their permanent dipoles. Stronger than dispersion forces.

Hydrogen Bonding

A particularly strong type of dipole-dipole interaction between molecules with N-H, O-H, or F-H bonds.

Polarizability

The degree to which the electron cloud of an atom or molecule can be distorted, influencing the strength of dispersion forces.

Ion-Dipole Interaction

Interaction between an ion (full charge) and the dipole (partial charge) of a molecule.

Hydrophobic Effect

The tendency of non-polar substances to aggregate in water, driven by water's preference to maintain H-bonding.

Mensicus

The curved surface of a liquid in a narrow tube, resulting from IMFs and attractions to the tube's surface.

Cohesive Forces

Intermolecular forces that cause a substance to stick to itself.

Adhesive Forces

Intermolecular forces that cause a substance to stick to other surfaces.

Gas

A state of matter characterized by disorder, no fixed volume, compressibility, and no fixed density.

Liquid

A state of matter characterized by disorder, fixed volume and density, and the ability to flow; it is non-compressible.

Solid

A state of matter characterized by order, fixed volume and density, and non-compressibility.

Vaporization

The phase transition from liquid to gas, requiring the input of heat to overcome IMFs.

Condensation

The phase transition from gas to liquid, releasing heat upon IMF formation.

Melting (Fusion)

The phase transition from solid to liquid.

Sublimation

The phase transition from solid to gas.

Deposition

The phase transition from gas to solid.

Heating Curve

A graph displaying the amount of heat needed to promote a phase transition.

Heat Capacity (Cs)

The amount of heat required to change the temperature of a substance, which varies for solids, liquids, and gases.

Phase Transition Heat

At a phase transition, heat goes entirely into overcoming or forming IMFs, not changing temperature.

Phase Diagram

A diagram showing the physical states of a substance under different conditions of temperature and pressure.

Vapor Pressure

The pressure exerted by a vapor in equilibrium with its condensed phases (solid or liquid) at a given temperature.

Triple Point

The temperature and pressure at which a substance can exist in equilibrium in three different phases (solid, liquid, and gas).

Critical Point

The point on a phase diagram at which the substance has properties intermediate between those of the liquid and gas phases.

Boiling Point

The temperature at which the vapor pressure of a liquid equals the surrounding environmental pressure.

Sublimation

The process by which a solid changes directly into a gas.

Deposition

The change of state from a gas to a solid.

Enthalpy of Vaporization

The heat required to vaporize one mole of a liquid at a specified temperature.

Solution

Homogeneous mixture of two or more substances.

Solvent

Substance in excess in a solution.

Solute

Substance not in excess in a solution.

Electrolyte

A substance that ionizes in H2O (dissociates into ions).

Strong Electrolyte

Completely ionizes in solution.

Weak Electrolyte

Partially ionizes in solution.

Non-Electrolyte

Does not ionize in solution.

Immiscible

Two substances that do not form a solution.

Colloid

A homogeneous mixture formed by ultramicroscopic particles in dispersal of a medium.

Molarity (M)

moles A / L solution

Molality (m)

moles A / kg solvent

Mole fraction (XA)

nA / n total

Henry's Law

C = k * PA, where C is the concentration of a gas, k is Henry's law constant, and PA is the partial pressure of the gas.

Colligative Properties

Properties of solutions that depend on the number of solute particles, not their identities.

Dissociation of Solutes

Solutes that dissociate produce more particles in solution than those that do not.

Van't Hoff factor (i)

The number of moles of dissolved particles produced from each mole of solute compound.

Effect of Increased Solute Particles

As the number of dissolved particles increases, freezing point and vapor pressure decreases, while boiling point and osmotic pressure increase.

Effect of Solute on Liquid Range

The presence of a solute expands the liquid range of the solvent.

Freezing Point Depression Equation

Tf = Tf° - ΔTf = Tf° - i * m * Kf, where Tf is the new freezing point, Tf° is the freezing point of the pure solvent, i is the van't Hoff factor, m is the molal concentration, and Kf is the freezing point depression constant.

Freezing Point Depression Change Equation

ΔTf = i * m * Kf, where ΔTf is the size of the change in freezing point, i is the van't Hoff factor, m is the molal concentration, and Kf is the freezing point depression constant.

Boiling Point Elevation Equation

Tb = Tb° + ΔTb = Tb° + i * m * Kb, where Tb is the new boiling point, Tb° is the boiling point of the pure solvent, i is the van't Hoff factor, m is the molal concentration, and Kb is the boiling point elevation constant.

Raoult's Law

The vapor pressure of a solution is equal to the mole fraction of the solvent times the vapor pressure of the pure solvent.

Raoult's Law Equation

Pi = xi * Pi°, where Pi is the vapor pressure of component i above the solution, xi is the mole fraction of i in the solution, and Pi° is the vapor pressure of pure component i under current conditions.

Volatile Compounds

Compounds with significant vapor pressures.

Nonvolatile Compounds

Compounds with negligible vapor pressures.

Raoult's Law

The vapor pressure of a solution is directly proportional to the mole fraction of the solvent in the solution: Psol'n = Xsolvent * P°solvent

Ideal Solution

A solution that follows Raoult's Law, where the interactions between solvent and solute are similar to those within the individual components.

Volatile Solute

If the solute is also volatile, the total vapor pressure of the solution is the sum of the partial pressures of each volatile component: Psol'n = XA * P°A + XB * P°B

Deviation from Raoult's Law (Positive)

Occurs when solvent-solute interactions are less favorable than solvent-solvent or solute-solute interactions, leading to a higher vapor pressure than predicted.

Deviation from Raoult's Law (Negative)

Occurs when solvent-solute interactions are more favorable than solvent-solvent or solute-solute interactions, leading to a lower vapor pressure than predicted.

Osmosis

The movement of a solvent through a semi-permeable membrane from a region of lower solute concentration to a region of higher solute concentration, driven by entropy.

Osmotic Pressure (π)

The pressure required to stop osmosis. Can be calculated using the formula π = iMRT, where i is the van't Hoff factor, M is molar concentration, R is the ideal gas constant, and T is temperature in Kelvin.

Reverse Osmosis

A process where pressure greater than the osmotic pressure is applied to a solution, forcing the solvent through a semi-permeable membrane, effectively separating it from the solutes.

Arrhenius Acid

A substance that produces hydrogen ions (H+) in water (really H3O+).

Arrhenius Base

A substance that produces hydroxide ions (OH-) in water.

Lewis Acid

An electron pair acceptor.

Lewis Base

An electron pair donor.