UPCAT SCI CHEM - COMPOUNDS 2

Metallic bond

Electrostatic attraction between metal cations and a sea of delocalized electrons. Stable because cations and electron sea mutually attract. Properties: conducts electricity, malleable, shiny.

Ionic bond

Complete transfer of electrons from metal to nonmetal. Metal = low ionization energy (gives), nonmetal = high electronegativity (takes). Opposite charges attract = the bond.

Covalent bond

Electrons SHARED between two nonmetals with similar electronegativities. Both nuclei attract shared electrons between them.

Coordinate covalent bond

One atom donates BOTH electrons to form a bond. Also called dative bond. Example: NH3 donates lone pair to H+ → NH4+

Compound vs molecule

Molecule = 2+ atoms bonded (can be same element, e.g. O2). Compound = 2+ DIFFERENT elements bonded (e.g. H2O). NaCl = compound but NOT molecule (ionic).

Binary covalent naming rules

Use prefixes (mono-deca) for atom count. Drop mono for first element. Second element gets -ide. Drop vowel before oxide. Example: N2O5 = dinitrogen pentoxide.

Type I ionic naming

Metal always ONE charge. Name = metal + nonmetal root + -ide. No Roman numerals, no prefixes. Example: NaCl = sodium chloride.

Type II ionic naming

Metal has VARIABLE charge. Work backwards from anion's charge to find metal's charge. Add Roman numeral. Example: FeCl3 → Cl=-1 x3=-3, Fe=+3 → iron(III) chloride.

Criss-cross method

Swap absolute values of charges as subscripts. Reduce to lowest terms. Use parentheses for polyatomic ions when subscript >1.

Oxidation number of free elements

Always 0. Not bonded, no electron transfer.

Oxidation number of H (normal)

+1 — less electronegative, gives electron away.

Oxidation number of H in metal hydrides

-1 — metal is less electronegative than H so H takes the electron. Examples: NaH, CaH2, LiH.

Oxidation number of O (normal)

-2 — very electronegative, pulls 2 electrons.

Oxidation number of O in peroxides

-1 — bonded to identical O, can't steal from each other, split equally. Example: H2O2.

Oxidation number of F

Always -1 — most electronegative, always takes electrons.

Oxidation number of halogens bonded to O

O is more electronegative than most halogens so O pulls electrons AWAY making halogen positive. Example: Cl in HClO = +1, Cl in ClO3- = +5. Exception: F is always -1 (F is more electronegative than O). 😊

Octet rule exceptions

1. Expanded octet: Period 3+ (P, S, Cl, Xe) use d orbitals, can hold 10+ electrons. 2. Incomplete octet: Be=4e, B=6e. 3. Duet rule: H max 2 electrons.

Percent composition formula

% = (n x molar mass of element / molar mass of compound) x 100

Formal charge formula

FC = V - N - B/2. V=valence electrons, N=nonbonding electrons, B/2=half bonding electrons. Sum of all FC = overall charge.

Formal charge circle shortcut

Draw circle around atom. Lone pairs count fully + 1 per bond tip. FC = valence electrons - electrons inside circle.

Linear shape

2 bonds, 0 lone pairs, 180°. Nonpolar. Example: CO2.

Trigonal planar shape

3 bonds, 0 lone pairs, 120°. Nonpolar. Example: BF3.

Tetrahedral shape

4 bonds, 0 lone pairs, 109.5°. Nonpolar. Example: CH4.

Trigonal pyramidal shape

3 bonds, 1 lone pair, 107°. Polar. Example: NH3.

Bent shape

2 bonds, 2 lone pairs, 104.5°. Polar. Example: H2O.

Electron pair vs molecular geometry

Electron pair = counts ALL pairs (bonds + lone pairs). Molecular = bonding pairs only. Example: NH3 = tetrahedral electron pair, trigonal pyramidal molecular.

Polar molecule

Central atom HAS lone pairs OR different surrounding atoms → dipoles don't cancel.

Nonpolar molecule

Central atom NO lone pairs + all surrounding atoms same → dipoles cancel. Example: CO2, CH4, BF3.

London Dispersion Force

Weakest IMF. Temporary/instantaneous dipoles. Present in ALL molecules. Only IMF for nonpolar molecules. Van der Waals.

Dipole-Dipole Force

IMF between polar molecules. Permanent dipoles attract. Van der Waals.

Van der Waals forces

Umbrella term for LDF + Dipole-Dipole only. Does NOT include H-bonding or ion-dipole.

Hydrogen Bonding

H bonded to F, O, or N attracts lone pairs on F, O, N of neighboring molecule. H2S also shows H-bonding.

Ion-Dipole Force

Strongest IMF. Full ion + polar molecule. Example: NaCl in water.

IMF strength order

London Dispersion < Dipole-Dipole < Hydrogen Bonding < Ion-Dipole

HIGH IMF properties

Boiling point HIGH, Melting point HIGH, Viscosity HIGH, Surface tension HIGH, Vapor pressure LOW, Heat capacity HIGH.

Binary acid naming

hydro + halogen root + ic acid. Only aqueous. Example: HCl(aq) = hydrochloric acid. Gas form = hydrogen chloride.

Oxyacid naming

-ate anion → -ic acid (H2SO4 = sulfuric acid). -ite anion → -ous acid (H2SO3 = sulfurous acid).