General Chemistry Notes

Development of a New Atomic Model

  • Rutherford's model could not explain color changes upon heating or chemical properties.
  • Bohr proposed electrons exist in specific circular paths (orbits).
    • Ground state: lowest energy state.
    • Excited state: higher energy state.
    • Energy released as electromagnetic radiation (colored light) when electrons drop to original orbital.
    • Elements emit specific line-emission spectra; electrons exist in specific energy states.

Electromagnetic Radiation

  • Wavelength (λ\lambda): distance between wave peaks (cm, nm, Å).
  • Frequency (ν\nu): number of peaks passing a point per second.
  • Electromagnetic radiation travels by waves, forming the electromagnetic spectrum.
  • Photoelectric effect: light has particle nature (photons).
  • Light exhibits both wave-like and particle-like properties.
  • Max Planck: E=hνE = h\nu, where hh (Planck’s constant) = 6.626x10−346.626 x 10^{-34} Js.

Quantum Mechanics

  • Louis de Broglie: Electrons behave like waves with specific frequencies.
  • Werner Heisenberg: it is impossible to determine both position and velocity of an electron.
  • Schrödinger: equations describe electron motion and energies.
  • Quantum mechanical model determines allowed electron energies and locations.
  • Schrödinger’s equations define atomic orbitals and electron properties (4 quantum numbers).

Quantum Numbers and Atomic Orbitals

  • Atomic orbital: region of high probability of finding an electron which can hold only 2 electrons.
  • Principal quantum number (n): main energy level; higher n means larger atom, electron farther from nucleus.
    • The principle quantum number is the same as the number of sublevels within that principle energy level.
  • Angular momentum quantum number (l): shape of orbital (sublevel).
    • Letters for l: s, p, d, f.
    • Each sublevel corresponds to a different orbital shape.
  • Magnetic quantum number (m): 3D orientation of orbital (x, y, z axes).
  • Spin quantum number (s): electron spin (CW or CCW).

Electron Configurations

  • Aufbau Principle: electrons occupy lowest energy orbitals first.
  • Pauli Exclusion Principle: no two electrons have the same four quantum numbers.
  • Hund’s Rule: orbitals of equal energy are each occupied by one electron before any are occupied by a second electron. All electrons in singly occupied orbitals have the same spin

Periodic Table and Electron Configurations

  • Period corresponds to the principal energy level (n).
  • Block corresponds to the sublevel being filled.
  • Exceptions: Filled sublevels are stable; half-filled sublevels are also relatively stable.

Periodic Law

  • Elements arranged by atomic number have similar properties in the same group.
  • Metals: left side of the table; conduct heat/electricity, lustrous, ductile, malleable.
  • Nonmetals: right side; insulators, dull, brittle.
  • Metalloids: staircase line; properties of both metals and nonmetals.

Groups

  • Metals: alkali (Group 1), alkaline earth (Group 2), transition (Groups 3-12), inner transition (lanthanides/actinides).
  • Nonmetals: chalcogens (Group 16), halogens (Group 17), noble gases (Group 18).
  • Hydrogen (1s11s^1) is unique.
  • Noble Gases have a full outer energy level (s and p).
  • s and p blocks are main group elements; d block is transition metals; f block is inner transition metals.

Factors Affecting Trends

  • Nuclear charge: more protons = more electron pull.
  • Shielding effect: inner electrons block outer electrons from nuclear pull.

Atomic Radius

  • Half the distance between nuclei of identical bonded atoms.
  • Decreases across a period (increasing nuclear charge).
  • Increases down a group (increasing shielding).

Ions and Valence Electrons

  • Ion: atom or group with a charge.
  • Valence electrons: outermost energy level electrons involved in bonding.
    • Group 1 & 2: equal to group number.
    • Groups 13-18: group number minus 10.

Ion Formation

  • Metals lose electrons to form positive ions (cations).
  • Nonmetals gain electrons to form negative ions (anions).

Ionization Energy

  • Energy to remove an electron.
  • Increases across a period (increasing nuclear charge).
  • Decreases down a group (increasing shielding).

Ionic Radii

  • Cations (+ ions) are smaller than their parent atoms because there are more protons attracting the remaining electrons.
  • Anions (- ions) are larger than their parent atoms because there are more electrons being attracted by the same number of protons.

Electron Affinity

  • Energy change when an electron is acquired.
  • Generally increases across periods (increasing nuclear charge).
  • Generally decreases down groups (shielding effect).

Electronegativity

  • Ability of an atom to attract electrons in a compound.
  • Metals: low electronegativity.
  • Nonmetals: high electronegativity.
  • Increases across a period.
  • Decreases down a group.
  • Fluorine is the most electronegative element.

Chemical Bonds

  • Force holding atoms together (electrical attraction).
  • Molecular compounds: covalent bonds (sharing electrons).
    • Nonmetals only.
    • Covalent compounds are also referred to as molecules
  • Ionic compounds: ions attracted to each other.
    • Metal and a nonmetal
    • Form formula units (simplest ratio of ions).
    • Are also referred to as salts
    • Charges on the ions must cancel each other out to equal zero

Bond Polarity

  • Determined by electronegativity difference between atoms.
    • Nonpolar covalent: 0 to 0.3.
    • Polar covalent: 0.3 to 1.7.
    • Ionic: 1.7 to 3.3.
  • Nonpolar covalent bond: electrons shared equally.
  • Polar covalent bond: electrons shared unequally (more electronegative atoms get a slightly negative charge δ−\delta^-).
  • Ionic bond: electrons transferred.

Electron Dot Diagrams

  • Symbol represents nucleus and inner electrons.
  • Use group number or electron configuration to find the number of outer shell electrons
  • Each side represents an orbital with only 2 electrons maximum
  • Add dots to represent the electrons

Covalent Bonding

  • Atoms gain, lose, or share electrons to achieve eight electrons in the highest occupied energy level (octet rule).
  • Exceptions exist.
  • Bond energy: energy to break a covalent bond (large bond energy = strong bond).

Multiple Bonds

  • Single covalent bond: sharing one pair of electrons.
    • Represented by a dash in structural formulas.
  • Unshared pair: pair of valence electrons not shared (lone pair).
  • Double covalent bond: sharing two pairs of electrons.
  • Triple covalent bond: sharing three pairs of electrons.
  • Coordinate covalent bond: both electrons come from the same atom.

Naming Binary Molecular Compounds

  • Write less electronegative atom first.
  • Add a prefix to each name only if more than one atom except for the first atom.
  • Second atom always has numerical prefix and -ide suffix.
  • Drop “o” or “a” on prefix if element starts with a vowel.

Chemical Formulas

  • Formula mass: sum of average atomic masses in atomic mass units (u).
  • Molar mass: numerically equal to formula mass in grams per mole (g/mol).
  • Subscripts indicate number of atoms/ions in a formula unit and number of moles in one mole of compound.

Molar Mass and Conversions

  • Molar mass converts mass to moles.
  • 1 mol = molar mass in grams.
  • Avogadro’s hypothesis: equal volumes of gases at same temperature and pressure have equal particles.
  • STP (standard temperature and pressure): 0°C and 1 atm.
  • At STP, 1 mole of any gas = 22.4 L.

Types of Chemical Reactions

  • Combination: A + X → AX.
    • Metals + oxygen → metal oxides (e.g., M2OM_2O for Group 1, MOMO for Group 2).
    • Metals + halogen → metal halides (e.g., MXMX for Group 1, MX2MX_2 for Group 2).
    • Active metal oxides + water → metal hydroxides.
  • Decomposition: AX → A + X.
    • Binary compounds break into elements.
    • Metal carbonates → metal oxide + carbon dioxide.
    • Metal hydroxides → metal oxide + water.
    • Metal chlorates → metal chloride + oxygen.
    • Oxyacids → water + nonmetal oxide.
  • Single Displacement: A + BX → B + AX.
    • Activity series predicts reaction occurrence (more reactive replaces less reactive).
  • Double Displacement: AX + BY → AY + BX.
    • Ions exchange in aqueous solution.
    • Reaction occurs if a precipitate, insoluble gas, or molecular compound forms.
  • Combustion: substance + oxygen → carbon dioxide + water (usually).

Chemical Reactions Described

  • Chemical reaction: substances transformed into different substances.
  • Chemical equation: formulas of reactants and products with symbols.
    • Reactants: original substances.
    • Products: resulting substances.
  • Indications: heat/light, gas, precipitate, color change, odor change.

Balancing Equations

  • Must represent known facts and contain correct formulas.\n* Law of conservation of mass must be satisfied.
  • Word equation: names of reactants and products.
  • Formula equation: chemical formulas without amounts.
  • Coefficients indicate relative amounts (ratio).
  • Must meet the Law of Conservation - can NOT create or destroy matter
  • May NOT change subscripts
  • May ONLY change coefficient
  • Balanced equation: same elements and number of atoms on both sides.
  • Balance polyatomic ions as a unit and balance H and O last.

Molecular Geometry (VSEPR)

  • Valence Shell Electron Pair Repulsion model predicts molecule geometry.
  • Electrons repel each other, determining shape.
  • Unshared pairs repel more than shared pairs.
  • Double/triple bonds count as one bond.

Molecular Shapes

  • Linear: 2 atoms (AB structure); 180° angle.
  • Trigonal Planar: 3 shared pairs, 0 unshared pairs (AB3 structure); 120° angle.
  • Tetrahedral: 4 shared pairs, 0 unshared pairs (AB4 structure); 109.5° angle.
  • Pyramidal: 3 shared pairs, 1 unshared pair (AB3E structure); 107° angle.
  • Bent: 2 shared pairs, 2 unshared pairs (AB2E2 structure); 104.5° angle.

Polar Molecules

  • One end is slightly negative, other is slightly positive (dipoles).
  • Symmetrical shape cancels bond polarity.

Intermolecular Forces

  • Weaker than ionic or covalent bonds.
  • Determine state (gas, liquid, solid).
  • London dispersion forces: between nonpolar molecules (induced dipole).
  • Dipole interactions: attractions between dipoles.
  • Hydrogen bonding: strong dipole when H bonds to electronegative atom.

Metallic Bonds

  • Vacant orbitals overlap, forming a “sea of electrons” (delocalized electrons).
  • Metallic bond: attraction of metal nuclei to the electron sea.
  • Electrons move freely, making metals conductive and shiny.
  • Atoms slide easily, making metals malleable and ductile.
  • Number of electrons determines bond strength.

Alloys

  • Mixtures of elements (at least one metal).
  • Superior properties.
  • Substitutional alloy: atoms of equal size replace each other.
  • Interstitial alloy: smaller atoms fit between larger atoms.

Naming Acids and Bases

  • Acids produce H+H^+ in water.
    • ide → hydro__ic acid
    • ate → __ic acid
    • ite → __ous acid
  • Bases produce OH−OH^- in water (name as ionic compound).