RJAV GenChem

$$\text{Solid} > \text{Liquid} > \text{Gas} > \text{Plasma}$$

Relative order of Intermolecular Force of Attraction (IFA) strength among the states of matter

$$\text{Plasma} > \text{Gas} > \text{Liquid} > \text{Solid}$$

Relative order of enthalpy (heat or reaction energy) among the states of matter

Recombination

Phase transition directly from plasma to gas, also known as deionization

Ionization

Phase transition directly from gas to plasma

Compound

Pure substance comprised of two or more chemically united elements that can only be separated via chemical means

Extrinsic (extensive) properties

Properties of matter that depend on the amount of matter present, such as length, mass, volume, pressure, entropy, and enthalpy

Intrinsic (intensive) properties

Properties of matter that are independent of the amount of matter present, such as density, viscosity, velocity, temperature, and color

Law of Conservation of Mass

Chemical law stated by Antoine Lavoisier proposing that mass is neither created nor destroyed in a chemical reaction

Law of Definite Proportions (Constant Proportions)

Chemical law stated by Joseph Proust proposing that chemical compounds always contain elements in fixed mass ratios

Law of Multiple Proportions (John Dalton)

Law stating that when two elements combine to form more than one compound, the masses of one element combining with a fixed mass of the other are in ratios of small whole numbers

Billiard Ball Model

Atomic model proposed by John Dalton describing matter as made of indivisible, indestructible spherical particles

Plum Pudding (Raisin Bread) Model

Atomic model proposed by J.J. Thomson describing electrons embedded within a positively charged framework

Oil Drop Experiment

Experiment performed by R.A. Millikan used to measure the accurate charge and mass of an electron

Neutron

Subatomic particle discovered by James Chadwick that carries no electrical charge

Quantum Mechanical Model

Atomic model developed by Erwin Schrödinger estimating the probability of finding an electron in an electron cloud or orbital

Eugene Goldstein

Scientist who discovered anode rays

Electrochemistry

Field of chemistry involving particle separation based on electron charge, such as capillary electrophoresis which separates compounds based on electrophoretic mobility

Anode

Positively charged electrode ($+\text{electrode}$) where oxidation occurs

Cathode

Negatively charged electrode ($-\text{electrode}$) where reduction occurs

RED CAT ELECT IN

Mnemonic standing for: REDuction happens in CAThode where ELECTrons get IN

VILEORA

Mnemonic standing for: Valence Increase, Loses $e^-$ (electrons), undergoes Oxidation, Reducing Agent

VDGEROA

Mnemonic standing for: Valence Decrease, Gains $e^-$ (electrons), undergoes Reduction, Oxidizing Agent (e.g., $KMnO<em>4$, $Na</em>2Cr<em>2O</em>7$)

Isotopes

Atoms of the same element with the same atomic number (protons) but differing atomic mass (neutrons), such as $^1H$, $^{12}C$, $^{14}N$, $^{32}S$, and $^{35}Cl$

Isobars

Atoms of different elements that have the same atomic mass but different atomic numbers

Isotones

Atoms that have the same number of neutrons but a different number of protons

Isomers

Compounds that share the same molecular formula but differ in structure

Molecule

An aggregate of two or more atoms in a definite arrangement held together by chemical bonds

Ions

Chemical species with a net positive ($+\text{charge}$) or negative ($-\text{charge}$) charge

Empirical formula

Chemical formula showing the simplest whole number ratio of atoms in a compound (e.g., $CH<em>2O$ for glucose $C</em>6H<em>{12}O</em>6$)

Intermolecular Forces (van der Waals / Electrostatic)

Weak and short-lived forces between molecules created by molecular polarizability when two uncharged atoms ($n^0$) approach closely

Keesom orientation force

Dipole-dipole interaction (D-D) between polar molecules, such as between water molecules

Debye induction force

Dipole-induced dipole interaction (D-ID) between a polar and a nonpolar molecule, such as between water and benzene

London dispersion force

Induced dipole-induced dipole interaction (ID-ID), which is the weakest intermolecular force present in all molecules, particularly aromatics like benzene-benzene

Hydrogen bonding

Strongest intermolecular force of attraction, occurring between hydrogen ($H$) and electronegative atoms ($S$, $O$, $N$, $X$)

Intramolecular Forces

Forces of attraction that exist within a single molecule, including covalent and ionic bonds

Covalent bond

Bond formed by the sharing of electrons between nonmetal atoms, such as glycosidic bonds (ether bond $S-O-S$) and peptide bonds (amide bond)

Ionic bond

Bond formed by the transfer of electrons from a metal to a nonmetal (e.g., $NaCl$)

Lone pair

Pair of valence electrons that are not shared with another atom in a covalent bond

Valence Shell Electron Pair Repulsion (VSEPR) theory

Theory predicting molecular geometry and unbonded electron pairs based on electron pairs arranging themselves as far apart as possible to prevent repulsion

Linear molecular geometry

Geometry with a bond angle of $180^{\circ}$, characteristic of $CO_2$ and alkynes ($Sp$ hybridization)

Trigonal planar molecular geometry

Geometry with a bond angle of $120^{\circ}$, characteristic of $BF_3$ and alkenes ($Sp^2$ hybridization)

Tetrahedral / Bent molecular geometry

Geometry with a bond angle of $109.5^{\circ}$, characteristic of $CCl<em>4$, $H</em>2O$, and alkanes ($Sp^3$ hybridization)

Trigonal bipyramidal molecular geometry

Geometry with bond angles of $90^{\circ}$ and $120^{\circ}$, characteristic of $PF_5$ (an exemption to the octet rule)

Octahedral molecular geometry

Geometry with all $90^{\circ}$ bond angles, characteristic of $SF_6$ (an exemption to the octet rule)

Valence bond theory

Theory stating that chemical bonds are formed by the sharing of electrons from overlapping atomic orbitals

Sigma bond ($\sigma$)

Stronger covalent bond formed by the head-on (headways) overlap of atomic orbitals, such as spherical $s$ orbitals

Pi bond ($\pi$)

Weaker covalent bond formed by the sideways overlap of atomic orbitals, such as dumbbell-shaped $p$ orbitals

Molecular orbital theory

Theory stating that chemical bonds are formed from the interaction of atomic orbitals to form molecular orbitals

Bonding molecular orbital

Lower energy, stable molecular orbital formed by the interaction of atomic orbitals (e.g., $\sigma$)

Antibonding molecular orbital

Higher energy, unstable molecular orbital formed by the interaction of atomic orbitals (e.g., $\sigma^*$)

Synthesis reaction

Reaction type also known as combination or direct union, represented by the general formula $A + B \rightarrow AB$

Decomposition reaction

Reaction type also known as analysis, represented by the general formula $AB \rightarrow A + B$

Complete combustion

Reaction of a hydrocarbon with abundant oxygen yielding carbon dioxide and water ($CH<em>4 + 2O</em>2 \rightarrow CO<em>2 + 2H</em>2O$)

Incomplete combustion

Reaction of a hydrocarbon with limited oxygen yielding carbon monoxide, solid carbon, and water ($CH<em>4 + O</em>2 \rightarrow CO + C<em>{(s)} + H</em>2O$)

Single displacement reaction

Chemical reaction represented by the general equation $AB + X \rightarrow AX + B$

Double displacement reaction

Reaction type also known as metathesis or exchange, represented by $AB + CD \rightarrow AC + BD$

Neutralization reaction

Double displacement reaction between an acid and a base forming a salt and water (e.g., $NaOH + HCl \rightarrow NaCl + H_2O$)

Precipitation reaction

Double displacement reaction resulting in an insoluble solid, such as $AgNO<em>3 + NaCl \rightarrow AgCl + NaNO</em>3$ which produces white curdy $AgCl$

Monovalent metal cations

Group 1 elements ($H$, $Li$, $Na$, $K$) and silver ($Ag$), which carry a charge of $+1$

Divalent metal cations

Group 2 elements ($Be$, $Mg$, $Ca$, $Sr$, $Ba$), zinc ($Zn$), and cadmium ($Cd$), which carry a charge of $+2$

Multivalent cations with $+1$ and $+2$ charges

Mercury ($Hg$) and copper ($Cu$)

Multivalent cation with $+1$ and $+3$ charges

Gold ($Au$)

Multivalent cations with $+2$ and $+3$ charges

Iron ($Fe$), cobalt ($Co$), and nickel ($Ni$)

Multivalent cations with $+3$ and $+5$ charges

Bismuth ($Bi$) and antimony ($Sb$)

Avogadro's number

Constant stating $1\text{ mole} = 6.022 \times 10^{23}\text{ atoms or molecules}$

Molarity ($M$)

Concentration formula defined as $M = \frac{n<em>{\text{solute}}}{L</em>{\text{soln}}}$

Molality ($m$)

Concentration formula defined as $m = \frac{n<em>{\text{solute}}}{kg</em>{\text{solvent}}}$

Normality ($N$)

Concentration formula calculated as $N = \frac{GEW}{L}$ or $N = M \times f$

Aufbau principle

Rule stating that electrons progressively occupy the sub-levels of lowest energy first

Subshell electron capacities

Maximum electron capacities per orbital subshell: $s=2$, $p=6$, $d=10$, and $f=14$

Pauli's exclusion theory

Principle stating that no two electrons in an atom can have the exact same set of four quantum numbers

Heisenberg's uncertainty theory

Principle stating that it is impossible to simultaneously and accurately determine both the position and velocity (momentum) of a particle

Hund's rule

Rule stating that orbitals of equal energy are filled singly before pairing up, and the most stable arrangement in subshells has the greatest number of parallel spins

Principal quantum number ($n$)

Quantum number ranging from $n = 1$ to $7$ that indicates the main energy level, orbital size, and distance of an electron from the nucleus

Azimuthal quantum number ($l$)

Quantum number ranging from $l = 0$ to $3$ that determines the angular momentum, shape of an orbital, and subshell type

Sharp subshell ($s$)

Subshell corresponding to $l = 0$ characterized by a spherical shape

Principal subshell ($p$)

Subshell corresponding to $l = 1$ characterized by a dumbbell shape

Diffuse subshell ($d$)

Subshell corresponding to $l = 2$ characterized by a clover leaf shape

Fundamental subshell ($f$)

Subshell corresponding to $l = 3$

Magnetic quantum number ($m_l$)

Quantum number ranging from $-l$ to $+l$ that specifies the orientation of an orbital in 3D space

Magnetic spin quantum number ($m_s$)

Quantum number with values of $+\frac{1}{2}$ (clockwise, incomplete spin) or $-\frac{1}{2}$ (counterclockwise, complete spin) indicating magnetic moment or rotation

Diamagnetism

Type of magnetism characterized by an atom or ion having no unpaired electrons

Paramagnetism

Type of magnetism characterized by an atom or ion having at least one unpaired electron

Boyle's law

Gas law stating that pressure and volume are inversely proportional at constant temperature, represented as $P<em>1 V</em>1 = P<em>2 V</em>2$ or $P \propto \frac{1}{V}$

Charles's law

Gas law stating that volume and absolute temperature are directly proportional at constant pressure, represented as $\frac{V<em>1}{T</em>1} = \frac{V<em>2}{T</em>2}$ or $V \propto T$

Gay-Lussac's law

Gas law stating that pressure and absolute temperature are directly proportional at constant volume, represented as $\frac{P<em>1}{T</em>1} = \frac{P<em>2}{T</em>2}$ or $P \propto T$

Combined gas law

Gas law combining pressure, volume, and temperature relationships, represented as $\frac{P<em>1 V</em>1}{T<em>1} = \frac{P</em>2 V<em>2}{T</em>2}$

Ideal gas law

Equation of state given by $PV = nRT$, where $R = 0.08205\,\text{L}\,\text{atm}\,\text{mol}^{-1}\,\text{K}^{-1}$

Standard temperature and pressure (STP)

Standard condition for gas calculations defined by $T = 273.15\,\text{K}$, $P = 1\,\text{atm} = 760\,\text{mmHg}$, and a molar volume of $V = 22.4\,\text{L}$

Avogadro's law

Gas law stating that equal volumes of different gases at the same temperature and pressure contain equal numbers of moles, represented as $\frac{V<em>1}{n</em>1} = \frac{V<em>2}{n</em>2}$ or $V \propto n$

Dalton's law of partial pressures

Gas law stating that the total pressure of a mixture of non-interacting gases equals the sum of the partial pressures of each gas, represented as $P<em>t = P</em>1 + P<em>2 + P</em>3$

Graham's law

Gas law stating that the rate of effusion or diffusion of a gas is inversely proportional to the square root of its density at constant temperature and pressure

Diffusion vs. effusion

Diffusion is the rate at which two gases mix, whereas effusion is the rate at which a gas escapes through a pinhole into a vacuum

Fick's 1st law

Law stating that the diffusion rate (flux) of a liquid or gas is directly proportional to the concentration gradient

Henry's law of gas solubility

Law stating that gas solubility is directly proportional to pressure, where decreasing temperature and increasing pressure dissolves more $CO_2$ in water

Van der Waals equation

Real gas equation $(P + \frac{a n^2}{V^2})(V - n b) = nRT$, where $a n^2$ accounts for internal pressure per mole and $n b$ accounts for incompressibility

Raoult's law

Law stating that $P<em>{\text{solution}} = X</em>{\text{solvent}} P_{\text{solvent}}$, where adding a non-volatile solute lowers the vapor pressure of the liquid

Absolute zero

Lowest possible temperature, defined as $0\,\text{K}$ (or $-273.15\,^{\circ}\text{C}$ / $-459.67\,^{\circ}\text{F}$)

Temperature conversion formulas

$^{\circ}\text{C} = \frac{^{\circ}\text{F} - 32}{1.8}$, $^{\circ}\text{F} = (^{\circ}\text{C} \times 1.8) + 32$, and $\text{K} = ^{\circ}\text{C} + 273.15$

Colligative properties

Properties of a solution that depend only on the amount of solute present, including vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure