Gen Chem Review

General Chemistry Rapid Review

Transition Metals

• Transition metals are the only elements that produce color.

• Color in transition metals is connected to their d-orbitals.

  • The d-orbital has an energy gap.

  • Energy Equations:

  • $E = hf$

  • $E = hc$

    • Where:

    • $h = 6.626 \times 10^{-34}$ (Planck's constant)

    • $c = 3.00 \times 10^8 \ m/s$ (speed of light)

• Color production involves electrons jumping between energy levels in the d-orbital.

Electron Configurations and Quantum Numbers

Example of Carbon (C)

• Electron configurations include:

  • Carbon has atomic number 6: $1s^2 2s^2 2p^2$

  • Follows Hund's Rule: every orbital in a sublevel is singly occupied before any orbital is doubly occupied.

• Paramagnetic vs. Diamagnetic:

  • Paramagnetic: contains unpaired electrons; shows magnetic properties.

  • Diamagnetic: all electrons are paired; does not exhibit magnetism.

Quantum Numbers

• Quantum number definitions:

  • Principal quantum number ($n$): indicates the main energy level.

  • Azimuthal quantum number ($l$): determines the shape of the orbital.

  • Magnetic quantum number ($m_l$): specifies the orientation of the orbital.

  • Spin quantum number ($m_s$): describes the intrinsic spin of the electron ($+\frac{1}{2}$ or $-\frac{1}{2}$).

• Possible values for $m_l$:

  • When $l = 2$, possible values are $-2, -1, 0, +1, +2$.

Ionic Radius Trends

• Ionic radii of species from noble gas configurations:

  • Trend: cations < neutral atoms < anions.

  • Example: Mg^{2+} < Na^{+} < F^{-} < O^{2-}

Electron Affinity and Electronegativity

• Electron affinity: the energy change when an electron is added to an atom.

• Electronegativity: the ability of an atom to attract electrons in a bond.

Acid-Base Concepts

• Brønsted-Lowry Theory:

  • Acid: proton donor

  • Base: proton acceptor

  • Example Reaction: (\text{HCl} + \text{H}2\text{O} \rightarrow \text{H}3\text{O}^+ + \text{Cl}^-)

• Coordinate Covalent Bond: a bond where both electrons are donated by one atom.

Intermolecular Forces

• Types of Forces:

  • Van der Waals/London dispersion forces: weak forces between nonpolar molecules.

  • Dipole-Dipole Forces: occur between polar molecules with dipoles.

  • Hydrogen Bonds: strong attractions between molecules containing H bonded to N, O, or F.

  • Ion-Ion forces: strongest attractions, e.g., between $Na^{+}$ and $Cl^{-}$.

Gibbs Free Energy and Thermodynamics

• Gibbs Free Energy Formula:

  • $\Delta G = -RT \ln K$

  • Where,

  • $R$ = universal gas constant,

  • $T$ = temperature in Kelvin,

  • $K$ = equilibrium constant

• Spontaneous vs. Non-Spontaneous Reactions:

  • Spontaneous: if \Delta G < 0.

  • Non-Spontaneous: if \Delta G > 0.

Redox Reactions

• Oxidation-Reduction Processes:

  • OIL RIG: Oxidation Is Loss, Reduction Is Gain.

  • Example Reaction: $\text{Zn} + \text{CuSO}<em>4 \rightarrow \text{ZnSO}</em>4 + \text{Cu}$

  • Oxidation: Loss of electrons (increases oxidation state).

  • Reduction: Gain of electrons (decreases oxidation state).

• Reduction Potentials:

  • Measurement of how strongly a species wants to be reduced; higher potentials indicate a stronger tendency to gain electrons.

Electrochemical Cells

• Galvanic Cells: spontaneous reactions that generate electrical energy.

• Electrolytic Cells: non-spontaneous reactions that require electrical energy to proceed.

  • Example: $E<em>{cell} = E</em>{reduce} - E_{oxidize}$

  • Standard cell potential calculation involves using standard reduction potentials.

Stoichiometry

• Mole Concept:

  • A mole (mol) is defined as $6.022 \times 10^{23}$ particles.

• Molar Mass Calculation: Mass of one mole of a substance (g/mol).

• Stoichiometric Relationships:

  • Use the coefficients in a balanced chemical equation to relate the amounts of substances.

Types of Chemical Reactions

• Combination: two or more reactants form a single product.

• Decomposition: a single reactant breaks down into multiple products.

• Combustion: reactant (hydrocarbon) burns in the presence of oxygen to produce CO2 and H2O.

• Single Replacement: one element displaces another in a compound.

• Double Replacement: two compounds exchange partners.

• Neutralization: an acid reacts with a base to form water and a salt.

Chemical Kinetics

• Rate-Determining Step: the slowest step in a reaction mechanism, which controls the overall rate of the reaction.

• Collision Theory: the rate of reaction increases with the concentration of reactants, temperature, and the presence of a catalyst.

• Reaction Rates:

  • Rate law expressions quantitatively describe how the rate depends on the concentration of reactants:

  • Example: $\text{Rate} = k[A]^m[B]^n$

Equilibrium

• Dynamic Equilibrium: the rate of forward reaction equals the rate of the reverse reaction, and concentrations remain constant over time.

• Equilibrium Constant: represented as $K$, determined from the concentrations of products and reactants:

  • $K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$ for the reaction (aA + bB \rightleftharpoons cC + dD)

Thermodynamics and Properties of Water

• First Law of Thermodynamics:

  • Energy cannot be created or destroyed, only transformed.

• States of Matter: solid, liquid, gas, and transitions between phases, including melting, freezing, sublimation, and deposition.

• Gibbs Free Energy Change:

  • $\Delta G = \Delta H - T\Delta S$ where $\Delta H$ is the change in enthalpy and $\Delta S$ is the change in entropy.

Solutions and Solubility Rules

• Solubility: a measure of how well a solute dissolves in a solvent

  • Soluble: all alkali metal salts, NH4+, NO3-, etc.

  • Insoluble: metal oxides and hydroxides generally are insoluble unless combined with alkali metals or certain cations.

Acids and Bases

• Brønsted-Lowry Definition: acids are proton donors, while bases are proton acceptors.

• Strong Acids include: HCl, H2SO4, HI, HBr, and HNO3.

• Exceptions in acidity: HF is not a strong acid due to its electronegativity and small size of F-.