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 ($\lambda$): distance between wave peaks (cm, nm, Å).
• Frequency ($\nu$): 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: $E = h\nu$, where $h$ (Planck’s constant) = $6.626 x 10^{-34}$ 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 ($1s^1$) 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 $\delta^-$).
• 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., $M_2O$ for Group 1, $MO$ for Group 2).
  • Metals + halogen → metal halides (e.g., $MX$ for Group 1, $MX_2$ 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 $H^+$ in water.
  • ide → hydro__ic acid
  • ate → __ic acid
  • ite → __ous acid
• Bases produce $OH^-$ in water (name as ionic compound).