unit 2: molecular and ionic compound structures

2.1 types of chemical bonds

electronegativity: a measure of the ability of an atom (or group of atoms) to attract shared electron (make bonds)

• increases as you move up and to the right on the periodic table

• doesn’t rlly apply to the noble gases bc they are so stable

• you compare electronegativity by look at periodic trends

non-polar: equal sharing of electrons(non-metal+non-metal)

polar: unequal sharing of electrons (non-metal+non-metal)

ionic: transferring of electrons (molecule becomes either more pos. or more neg.)(metal+non-metal)

• the farther apart two elements are, the more ionic their molecule is

metallic: bonds between metals (called an alloy)

• contain a delocalized sea of valence electrons (free-moving electrons)

• metallic character increases as you go down and left on the period table

2.2 intramolecular force and potential energy

covalent bonds:

• bonds between two non-metals

• can be polar or non-polar

• can be single, double, or triple bonds

  • averaged together=resonance structures

potential/bond energy graphs:

• covalent bonds

• ionic bond equivalent= lattice energy graph

• the trend:

  • smaller atoms are more towards the beginning of the graph because they are closer together than larger atoms

  • molecules with more bond (triple/double) have deeper graphs because they have greater bond energy

lattice energy: the amount of energy needed to separate an ionic compound

• smaller radi→ more attraction→ more lattice energy

• closer together → share more electrons

2.3 structure of ionic solids

ionic solids:

• non-volatile

• high melting points

• do not conduct electricity

  • only conduct once melted or dissolved

    • electrons become free-moving→ electricity is conducted

• many are soluble in polar solvents and non-soluble in non-polar solvents