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Chapter 9

CHM 111 Chapter Notes - Chapter 9: Valence Electron, Ionic Compound, Lattice Energy


Department
Chemistry
Course Code
CHM 111
Professor
Lachgar
Chapter
9

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CHAPTER 9
Maximum stability occurs when an atom is isoelectronic with a noble gas
Lewis Dot Symbols: symbol of element & 1 dot per valence electron
o Don’t really work for transition metals due to incomplete lower subshells
Ionic bonds: electrostatic forces that hold together ionic compounds
o Low ionization energy = cations, High electron affinity = anions
Lattice energy of ionic compounds is the energy required to completely separate 1 mole of a
solid ionic compound into gaseous ions
o Measures stability of ionic solids, always positive - greater LE → higher melting
point
Born-Haber Cycle: relates lattice energies of ionics to IEs, EAs, and other properties
Coulomb’s Law
o q refers to charges
Covalent Bonding: sharing of electrons between 2 atoms
o Covalent compounds contain only covalent bonds
Nonbonding electrons = lone pairs
Lewis Structures show covalent bonding where shared electron pairs are shown as lines
Octet Rule: bonds formed until atom has 8 electrons - except H
Single bond: 1 e- pair shared, double bond has 2e- paired, triple has 3, etc.
Multiple bonds are shorter than single bonds and are more stable
Bond length: distance b/w nuclei of 2 covalently bonded atoms in a molecule
2 types of attractive forces in covalents: 1) bond enthalpy 2) intermolecular forces
IMFs are usually weak → reason why covalents are often gases/liquids/low melting solids
Electrostatic forces are strong in ionics
Electronegativity: relative measure of atom’s ability to attract toward itself the e- in bonds
o Pauling developed it, F= 4.0
o Low electronegativity = low IE = low EA, increases up periods and across groups
Large difference in electronegativity → ionic bonds (>2.0 difference)
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