vid 10 and 11 ionic and covalent bonds

lewis theory states that valence electrons are important in

determining chemical bonding

lewis theory states bonding happens so that

each atom attains noble gas configuration

ionic bond

complete transfer of electrons between metals and nonmetals because of a large difference in tendency to gain or lose e-

covalent bond

sharing of electrons between nonmetals

metallic bond

electron pools of delocalized electrons form between metals

enthalpy/∆H

the heat absorbed or released under constant pressure in a reaction

when ∆H > 0, reaction is

endothermic

when ∆H < 0, reaction is

exothermic

steps in formation of lattice

formation of gaseous metal atoms, cations, nonmetal atoms, nonmetal anions — large absorption of heat

lattice energy

that which is required to destroy the lattice - opposite sign but same magnitude of what was required to form the solid

lattice energy =

q1q2/rˆ2 → q are the anionic and cationic charges, r is the distance between them

down a group, lattice energy

decreases

across a period, lattice energy 

increass

ionic solids characteristics

hard, rigid, and brittle; high melting point; conduct electricity in liquid or molten state

phases of covalent interaction

as atoms get closer, there is attraction and PE decreases from zero

reaches an optimal distance and a PE minimum

too close and there is repulsion, PE increases

from potential energy vs internuclear distance graph we can obtain

bond length (x axis) and bond energy magnitude (y axis)

as number of bonds increase, bond length _ and energy _

decreases; increases

bond breakage is

endothermic

bond formation is

exothermic

when given the weird reactions and energy thing

∆H of overall reaction = last row = enthalpy of atomization (solid to gas) + ionization energy + bond energy + electron affinity - lattice energy