GenChem 1

Fusion, Liquefaction, Thawing

Melting (Solid to Liquid)

Recombination, Deionization

Plasma to Gas

Ionization

Gas to Plasma

Extrinsic Property

Dependent properties like length, mass/weight, volume, pressure, entropy, enthalpy, electrical resistance.

Intrinsic Property

Independent properties like density (SpGr), viscosity, velocity, temperature, color.

Law of Conservation of Mass/Matter

Antoine Lavoisier's principle that mass is always constant; matter cannot be created or destroyed.

Law of Definite/Constant Proportions

Joseph Proust's law stating that a given chemical compound always contains its component elements in fixed ratio by mass.

Law of Multiple Proportion – John Dalton

When 2 elements form more than 1 compound, it can be expressed in a fixed whole number (by mass). Ex: CO → 28g/mole, CO₂ → 44g/mole; C = 12g/mole, O = 16g/mole.

Law of combining weights

Proportions by weight when a chemical reaction takes place can be expressed in small integral units. Ex: MgO → 40g/mole (100%); Mg = 24g/mole (60%); O = 16g/mole (40%).

Democritus

Atomos - 'Indivisible'.

John Dalton

Billiard ball; proposed that matter is made up of atoms.

Dalton's Postulates

1. Elements are composed of indivisible, indestructible atoms. 2. Atoms alike for a given element (isotopes). 3. Atoms of different elements differ in size, mass & other properties (isobars). 4. Compounds are formed from 2 or more atoms at different elements. 5. Atoms combined in simple numerical ratios to form compounds.

J.J. Thompson

Proposed the Plum Pudding / Raisin bread model; e- in (+) framework.

Ernest Rutherford

Discoverer of proton; Nuclear model from Gold foil / α-scattering experiment; atom is mostly empty, (+) particles in nucleus.

Neil Bohr

Planetary model; mostly used.

Erwin Schrodinger

Quantum/Mechanical model; estimates the probability of finding an e- in certain position (i.e

Proton

(+) ion; Atomic number (basis of electronic configuration); Ernest Rutherford.

Electrons

(-) ion; p+ in uncharged state; negligible weight 1,836x lighter than p+; J.J. Thompson proposed the cathode ray tube; R.A. Millikan measured accurate charge and mass of e-.

Neutrons

No charge; Atomic mass (Nucleon) = p+ + n0; James Chadwick.

p+

Atomic # = 11.

n0

Atomic mass - p+ = 23-11 = 12.

e-

p+ in uncharged state: 11-1 = 10.

P

p = 15.

n

n = 16.

e

e = 18.

Eugene Goldstein

Discovered anode rays.

Electrochemistry

Particle separation based on e⁻.

Anode

Electrode with a positive charge; undergoes oxidation.

Cathode

Electrode with a negative charge; undergoes reduction.

RED CAT ELECT IN

Mnemonic: Reduction happens in cathode where electrons get in.

VILEORA

Valence Increase, Loses e⁻; undergoes oxidation; reducing agent.

VDGEROA

Valence Decrease, Gains e⁻; undergoes reduction; oxidizing agent.

Isotopes

Same p⁺/atomic number/element but differ in atomic mass.

Non-isotopes

e.g., ¹⁹F, ¹²¹Sn, etc.

Isobars

Same atomic mass, differ in elements.

Isomers

Same molecular formula, differ

• H Bonding

• Keesom

• Debye

• London

• Covalent

• Ionic

• Ether

• Amide

Lone pair

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

Valence bond theory

States that bonds are formed by sharing of electron from overlapping atomic orbitals (covalent)

Molecular bond theory

States that bonds are formed from interaction of atomic orbitals from molecular orbitals

Synthesis/ Combination/ Direct Union

A + B → AB

Decomposition/ Analysis

AB → A + B

Complete combustion example

CH₄ + O₂ → CO₂ + H₂O

Incomplete combustion example

CH₄ + O₂ → CO + C(s) + H₂O

Single Displacement

AB + X → AX + B

Double Displacement/ Metathesis/ Exchange

AB + CD → AC + BD

• Li CA ba CA na mg Al MN zn Cr fe co CD ni sn PB H2 cu ag hg PT au

• F cl br I

• Hg cu

• Au

• Fe co ni

• Bi SB

Recall

Answer recall

1 mole = 6.022 x 1023 atoms/ molecules

Avogadro's number

Aufbau principle

Atoms may be built by progressive filling of energy of main energy sub level (i.e., levels of lower energy levels are occupied first

• 2

• 10

• 18

• 36

• 54

• 86

Principal Quantum Number (n)

Main energy level; size of orbital (electron cloud), distance of e⁻ from nucleus. Example: O₂ is 1s² 2s² 2p⁴.

Azimuthal/ Angular Momentum Quantum Number (ℓ)

Angular momentum & shape of orbital; subshell. Values: ℓ = 0 (s: sharp), 1 (p: principle), 2 (d: diffuse), 3 (f: fundamental). Example: O₂ = ℓ = 1.

Magnetic Quantum Number (mℓ)

Orientation of orbital in space. Example: O₂ has mℓ = -1, 0, +1.

Magnetic Spin Quantum Number (ms)

Magnetic moment/Rotation. Values: +½ (clockwise), -½ (counterclockwise). Example: Oxygen has ms = +½.

Pauli’s exclusion theory

No 2 e⁻ will have the same set of quantum numbers ('exclusive').

Heisenberg’s uncertainty theory

Impossible to predict/accurately determine the particle’s velocity (position & momentum).

Hund’s rule

Orbitals are filled up singly before pairing up; most stable arrangement of e⁻ in subshells is the one with greatest no. of parallel spins.

• Boyle's/Mariotte

• Charle’s

• Gay-lussac

• Combined gas law

• Ideal gas law

• Avogadro's law

Dalton's Law of Partial Pressures

Total pressure in a mixture (non-interacting gases) is equal to the sum of the partial pressures of each gas. $P<em>t = P</em>1 + P<em>2 + P</em>3$.

Graham's Law

Rate of effusion (diffusion) and speed gas are inversely proportional to the square root of their density, given constant temperature and pressure.

Diffusion

Rate at which 2 gases mix.

Effusion

Rate at which gas escapes through a pinhole vacuum.

Fick's 1st Law

Diffusion rate (flux) of liquid or gas is directly proportional to the concentration gradient (from high concentration to low concentration).

Henry's Law of Gas Solubility

Pressure is directly proportional to solubility; decrease in temperature and increase in pressure (e.g., sealed container) allows more CO₂ to be dissolved in water.

Raoult's law

vapor pressure of a solution is dependent on the amount of nonvolatile solute added to solution

Open System

Allows exchange of energy and matter.

Closed System

Allows exchange of energy but not matter.

Isolated System

Does not allow exchange of both energy and matter; referred to as 'Adiabatic Walls'.

Surrounding

Everything outside the system.

State Function

Independent; depends only on initial and final states of the system (e.g., Enthalpy (H), Internal energy (U), Gibbs Free Energy (G), Entropy (S)).

Non-State Function

Dependent; includes Work and Heat.

Zeroth Law

If two systems are in thermal equilibrium respectively with a third system, they must be in thermal equilibrium with each other (a=c, b=c, a=b).

1st LAW: Law of conservation of Energy

Energy is neither created nor destroyed but can be transformed from one form to another.

Enthalpy (H)

H = U, P, V.

Hess' Law

ΔH is independent of reaction/steps that occurred (only the initial and final steps is the basis).

2nd LAW: Law of Entropy

No way but UP. For an isolated system, Total entropy can never decrease

Entropy (ΔS)

Measure of a system's thermal energy per unit temperature; degree of disorderliness or randomness. ΔS = (+) → spontaneous; increase (irreversible) – real case; ΔS = (−) → non-spontaneous; constant (reversible) – ideal case (in a steady state/equilibrium); ΔH → does not predict spontaneity.

3rd LAW

If an object reaches absolute zero temperature (0 K = -273.15 = -459.67 °F), the entropy of a perfect, solid, crystalline substance is zero at absolute 0 temperature.

Reaction Rate

The speed at which a chemical reaction occurs; typically measured by the change in concentration of reactants or products over time.

Rate Law

An equation that relates the reaction rate to the concentrations of the reactants, usually in the form of rate = k [A]^m [B]^n, where k is the rate constant and m and n are the reaction orders.

Collision Theory

Rate of chemical reaction is proportional to the number of collisions per time. Requirements for effective collision: Proper orientation, Activation energy (Ea) – minimum amount of energy required to initiate chemical reaction.

Transition Theory

Formation of Intermediate Complex; rate depends on Ea required to form an intermediate state (where new bonds are formed and old bonds are broken).

Le chatelier’s principle

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

• Sour, bitter

• Red, blue

• H2

• CO2

• Soap

• NaOH

• KOH

Recall

Answer recall

Common Ion Effect

Addition of a compound having an ion in common with the dissolved substance will result in: • Equilibrium shift (either to the left or right) • Suppressed ionization of the dissolved substance (WA or WB) • pH change.

Henderson-Hasselbalch/Buffer Pair Equation

For buffer solutions (WA + CB or WB + CA), the equations are: Weak acids: $pH = pK<em>a + ext{log} \frac{[ ext{salt}]}{[ ext{acid}]}$ Weak bases: $pH = pK</em>b + ext{log} \frac{[ ext{base}]}{[ ext{salt}]}$ *pH = pKa (@ half neutralization point).

• Unsaturated

• Saturated

• Supersaturated

Noyes Whitney equation

Dissolution rate is directly proportional to the solute surface area, solute concentration at boundary layer, and diffusion coefficient

• Basic, acidic

• Reducing, oxidizing

• Brittle

• Iodine

• Mercury

Antoine Lavoisier

First extensive list of elements (~33), distinguishing between metals and nonmetals.

Ionization Energy

Energy needed to remove the outermost electron in a neutral atom; increases across a period (→) and decreases down a group (↓).

Electron Affinity

Energy given when a neutral atom gains an extra electron (making it more negative); increases across a period (→) and decreases down a group (↓).

Electronegativity

Ability of an atom to attract an electron pair to itself during covalent bond formation; increases across a period (→) and decreases down a group (↓).

Atomic Radius

Half the distance between nuclei of 2 atoms; decreases across a period (→) and increases down a group (↓).

Johan Wolfgang Dobereiner

'Law of Triads'.

John Newlands

'Law of Octaves'; considered periods.

Dmitri Mendeleev

'Father of Modern Periodic Table'; Atomic Mass/Weight.