Chemistry CLEP

John Dalton

Proposed that matter is composed of atoms; these atoms have different identities called elements, which combine to form compounds; measured masses of reactants and products.

J.J. Thompson

Observed deflection of particles in a cathode ray tube; proposed that atoms are composed of positive and negative charges; developed the plum pudding model of the atom

Robert Millikan

Calculated the charge-to-mass ratio of electrons using oil drops falling in an electric field; surmised the charge of a single electron

Ernest Rutherford

Used the deflection of alpha particles in a cathode ray tube to discover that most of the atom is empty space, with protons and neutrons centered in the nucleus.

Niels Bohr

Determined that electrons exist around the nucleus at a fixed radius; electrons with higher energy exist farther from the nucleus. Electrons give off electromagnetic radiation when moving between energy levels.

Max Planck

Determined that energy is quantized, or composed of discrete bundles.

6.63 x 10^-34 J*sec

Planck's Constant (h)

3.00 x 10^8 m/sec

Speed of Light (c)

E \= hv

Energy of a Photon Formula (1)

E \= hc / wavelength

Energy of a Photon Formula (2)

Louis DeBroglie

Combined Einstein's relationship between mass and energy and the relationship between velocity and the wavelength of light. All particles with momentum have a corresponding wave nature.

Wavelength \= h / mv

Wavelength of Particles Formula

Heisenberg's Uncertainty Principle

It is impossible to simultaneously know the position and momentum of an electron.

Erwin Schrodinger

Attributed a wave function to electrons, describing the probability of where an electron might exist.

Orbitals

Regions of high probability where electrons might exist; broken into four levels: s, p, d, or f

Atomic Mass

The cumulative mass of all the particles in the atom; found by adding the masses of the protons and neutrons.

Units: Atomic Mass Units

AMU

Example for AMU

helium = 2 protons + 2 neutrons

= 4 AMU

Atomic Number

The number of protons in the nucleus of an atom, or the total nuclear charge. Also the number of electrons surrounding the nucleus.

Isotopes

Atoms with the same number of protons but different numbers of neutrons.

Example: carbon-12 (6 neutrons) vs carbon-14 (8 neutrons)

Isotopes

Atomic Weight

Molar mass of the element, or the mass in grams of one mole of atoms

Pauli Exclusion Principle

No two electrons can occupy the exact same energy level or have the same set of four quantum numbers

Quantum numbers

1. Principal (n)

2. Angular Momentum (l)

3. Magnetic (ml)

4. Magnetic Spin (ms)

Principal Quantum Number

The shell or energy level an electron occupies; values from 1-7. Electrons with higher values are farther from the nucleus.

Angular Momentum Quantum Number

The subshell the electron occupies; describes the shape of an electron's orbital.

n = 1: l = 0

(s)

n = 2:

l = 0 (s), 1 (p)

n = 3:

l = 0 (s), 1 (p), 2 (d)

n = 4:

l = 0 (s), 1 (p), 2 (d), 3 (f)

Magnetic Quantum Number

Represents the orbital position.

l = 0: ml = 0 (1 possible s orbital)

l = 1: ml = -1, 0, 1 (3 possible p orbitals)

l = 2: ml = -2, -1, 0, 1, 2 (5 possible d orbitals)

l = 3: ml = -3, -2, -1, 0, 1, 2, 3 (7 possible f orbitals)

Magnetic Spin Quantum Number

Each orbital contains at most 2 electrons: one with a positive spin (+1/2) and one with a negative spin (-1/2)

Dimagnetic

Elements that have paired electrons in each orbital; all subshells are filled. These elements aren't affected by magnetic fields

Paramagnetic

Elements that have an unpaired electron in at least one orbital; creates a magnetic field in the atom that responds to external magnetic fields.

Electron Configurations

Identify the number of electrons in each type of orbital at each energy level

Orbital notation

Identifies where each electron exists in each orbital.

Aufbau Principle

Electrons exist first at the lowest possible energy level, unless energy has put them into an excited state

Hund's Rule

Electrons enter orbitals of equal energy singly, with the same spin, before becoming paired.

Atomic radius

Decreases across a period and increases across a group; cations are smaller than neutral ions; anions are larger than neutral ions.

Ionization energy

The minimum amount of energy required to remove an electron from an atom; increases across a period and decreases down a group

Electronegativity

The tendency of an atom to attract electrons; increases across a period and decreases down a group

Electron affinity

The amount of energy released when an electron is added to a neutral atom; increases along a period and decreases down a group.

Covalent Bond

Occurs when pairs of electrons are shared by atoms. Atoms will bond with other atoms in order to gain more stability, which is gained by forming a full electron shell. Nonmetals will readily form these bonds with other nonmetals in order to obtain stability.

Valence electrons

The electrons in the outermost shell of an atom

Sigma bond

The first covalent bond between nonmetals

Pi bond

Each additional covalent bond between non-metals after a sigma bond; much weaker than sigma bonds.

Nonpolar covalent bond

Created when atoms share their electrons equally; usually occurs when two atoms have similar or the same electron affinity. The closer the values of their electron affinity, the stronger the attraction; normally the difference in electronegativity is \>0.4.

Polar covalent bond

Occurs when the electrons between atoms are not equally shared. The atom with the higher electronegativity will have a stronger pull for electrons, resulting in the molecule having a slightly positive side and a slightly negative side. Normally the difference in electronegativity is between 0.4 and 1.7.

Dipole moment

The measure of net molecular polarity. The larger the difference in electronegativities of bonded atoms, the larger the moment.

Network covalent structure

A chemical structure in which the atoms are bonded by a group of covalent bonds in a continuous network.

Lewis Structure

A model of the valence electrons that are involved in covalent bonding

Resonance structure

An attempt to model delocalized electrons

Hybridization

The process by which electrons mix traits of different atomic orbitals to create bonding orbitals; or, the concept of mixing atomic orbitals into new hybrid orbitals (with different energies, shapes, etc., than the component atomic orbitals) suitable for the pairing of electrons to form chemical bonds

VSEPR

Valence shell electron-pair repulsion; electron pairs will repel each other, making each electron pair as far away as possible from every other electron pair.

Expanded octets

Created when highly electronegative atoms bond to large central atoms and there is no space to allow either five or six electron pairs around the central atom. Require d-electrons to participate in hybridization.

Isomers

Molecules that have the same formula but different structure or arrangement of atoms.

Van der Waals Forces

An attraction of intermolecular forces between molecules created by the chance movement of electrons in a system of bonded atoms

Hydrogen bond

A hydrogen atom is involved with a polar intermolecular attraction to a more electronegative atom.

Polar intermolecular attraction

The slightly positive end of one molecule forms an electrostatic attraction to the slightly negative end of another molecule.

Alpha decay

An atomic nucleus emits an alpha particle, equal to the nucleus of a helium atom

Beta decay

An atomic nucleus emits an electron or positron, converting a neutron into a proton

Positron decay

A nucleus emits a particle that degrades into a positron; converts a proton into a neutron and a positron.

Gamma radiation

Given off in combination with alpha and beta decay; the rays given off are photons.

Half-life

The time it takes for 50% of an isotope to decay.

T \= 0.693 / k

Half-life formula

P1V1 = P2V2 (The volume of a gas is inversely proportional to its pressure, when temperature is constant)

Boyle's Law

V1T2 = V2T1 (The volume of a gas is directly proportional to temperature, when pressure is constant)

Charles's Law

P1T2 = P2T1 (The pressure of a gas is directly proportional to temperature, when volume is constant)

Law of Gay-Lussac

Ptotal = P1 + P2 + ... + Pn (The total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the gases in the mixture.)

Dalton's Law

V1n2 = V2n1 (The volume of a gas is proportional to the number of moles of gas present when temperature is constant)

Avogadro's Law

PV = nRT (Pressure x Volume = number of gas moles x ideal gas constant x absolute temperature)

Ideal gas law formula

0.082 L*atm / K*mol

Ideal gas constant (R)

273 K; 1.0 atm; 1.0 mol gas = 22.4 L gas

Standard Temperature and Pressure (3 components)

P(mm) = dRT (Pressure x molar mass = density x ideal gas constant x absolute temperature)

Ideal gas law formula (in terms of density)

1/2 (ma x va^2) = 1/2 (mb x vb^2)(Two gases at the same temperature and pressure will have the same kinetic energy; v = velocity; m = mass)

Graham's Law

ra^2 / rb^2 = Mb / Ma (gas molecules of smaller molar mass move faster than gas molecules of larger molar mass)

Graham's law of effusion

\[P + (an^2 / V^2)\](V - nb) = nRT (Pressure + number of moles squared x intermolecular attraction constant / volume of gas squared, x Volume - no. moles x space occupied by one mole, = number of moles x ideal gas constant x temperature)

Van der Waals Equation for real gases

Phase diagram

Shows the state of a substance at any given temperature and pressure

Critical point

The temperature and pressure point on a phase diagram above which the substance must exist as a gas.

Triple point

The temperature and pressure point on a phase diagram at which a substance may exist in all three phases

Vapor pressure curve

Defines the boundary between the liquid & gas phases on a phase diagram

Crystalline solids

Composed of structural units bounded by a specific geometric pattern. Example: table salt

Unit cell

The smallest repeating unit in a crystalline solid

Simple cubic unit cells

These have one atom at each of the corners of the cube; containing a total of one atom per unit cell.

Face-centered crystal

A simple cubic unit cell with one additional atom shared between two unit cells on each face of the cube; a total of three atoms per unit cell.

Body-centered crystal

A simple cubic unit cell with one additional atom in the center of the cube, for a total of two atoms per unit cell

Amorphous solids

These don't display a specific geometry; example: glass

Solvation

The interaction of solvent molecules with solute molecules to form loosely bonded combinations.

Hydration

The solvation process when water is the solvent.

Miscible solutions

These occur when one substance is soluble in all proportions with another substance.

Saturation

A solid solute is in equilibrium with dissolved solute.

Solubility

The molar concentration of dissolved solute at saturation

Supersaturation

A solution that contains more solute than required for saturation

P = kC (The amount of gas that can dissolve in a liquid is directly proportional to the partial pressure of the gas above the liquid.)

Henry's Law

M \= moles solute / liters solution

Molarity formula

pH \= - log [H+]

pH formula

Molality \= Moles solute / kilograms solvent

Molality formula

mole fraction \= moles solute / total solution moles

Mole Fraction formula

P = XP0 (vapor pressure of solution = vapor pressure of solvent x mole fraction of solvent)

Raoult's Law

∆T = (kb)(m)(i) (change in solvent boiling point = molal boiling point constant of solvent x molality of solute x van't hoff factor of solute)

Boiling point elevation; or, change in solvent boiling point