Chemistry Unit 7: Ionic Bonds + Dipoles & Covalent + Bond Enthalpy + Lewis Structures + VSPER Theory + Intermolecular Forces

bond energy

energy required to break a bond

bond length

distance where the energy is minimal

Coulomb’s law

governs attractions between charges; larger charges = more attraction, larger distance = weaker attraction

ion size

parents (protons) vs children (electrons)

lattice energy

amount of energy needed to combine individual gaseous ions into an ionic solid (how attracted ions are to each other)

determining lattice energy

more charges = higher lattice energy; THEN larger ion = lower lattice energy

electronegativity

ability of an atom to attract an electron (F)

polar

when one atom is more electronegative than another (unequal sharing of electrons)

nonpolar

equal sharing of electrons

dipolar and dipole moment

polar

lone pairs on central atom

polar

bond enthalpy

energy required to form or break bonds

bond enthalpy formula

deltaH = H_{bonds broken} - H_{bonds formed}

drawing lewis structure steps

sum up all valence electrons for total electrons in structure

determine central atom

draw bonds

add lone pairs

adjust for total electrons

H needs

one bond

C needs

4 bonds

N wants

3 bonds

oxygen wants

2 bonds

B is happy seeing how many valence

6 valence electrons

when can you break the octet

3rd period and below

P can see how many electrons

10 electrons

S can see how many electrons

12 electrons

resonance

when there is more than one valid lewis structure, molecule is avg of these structures

formal charge

VE brought - e^- owned in molecule

VSEPR

valence shell electron pair repulsion

structure of a molecule is determined by

electrons repelling each other

electron pair arrangement/geometry

how all electrons are laid out

molecular geometry

how only bonds are laid out

hybridization

mixing of orbitals to create “hybrid” orbitals (count # of bonds)

sigma bond

1st bond in every bonded atom

pi bond

all other bonds that are not the first

intermolecular forces

interaction between molecules

london dispersion forces

electrons temporarily move to charge one side of a molecule; more e^- = higher LDF; weakest IMF

dipole-dipole forces

polar molecules line up so + ends are attracted to - ends of another molecule

hydrogen bonding

special dipole-dipole bonding; H bonded to F, O, or N; strongest IMF

solubility

like dissolves like; polar dissolves polar, nonpolar dissolves nonpolar

hydrophillic

water loving polar substances

hydrophobic

water fearing nonpolar substances