Honors Chemistry Chapter 8 Study Guide

Iodine heptafloride

________ (IF7) can be written so that sulfur has twelve valence electrons.

Vander Waals

________ Forces: attractions between the molecules that are the two weakest attractions between molecules.

VSEPR

According to the ________ theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valence- electron pairs stay as far apart as possible.

Covalent Bond

________: joins two atoms together that are sharing a pair of electrons.

Nitrogen dioxide

________ (NO2) has an odd number of valence electrons and therefore does not satisfy the octet rule.

temperature

They have low in ________ melting points and boiling points.

Nonpolar

________ and caused by temporary asymmetrical dispersion of electrons around it.

Octet Rule

The ________: electron sharing usually occurs so that the atoms attain the electron configurations of noble gases (each having eight valence electrons)

Polar covalent bond

________: the electrons are shared unequally.

Formula Unit

________: the base of an ionic compound NaCl is an ionic compound so the ________ are Na and Cl.

Atoms

________ form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons.

Phosphorus pentachloride

________ (PCl5) can be written so that phosphorus has ten valence electrons.

Dipole Dipole interactions

________: when polar molecules are attracted to each other.

Boron Trifluoride

________ (BF3) can be written so that boron only has 6 valence electrons.

Dipole

________: in a polar molecule, one end of the molecule is slightly negative and the other slightly more positive.

Molecule

________: a neutral group of atoms joined together by covalent bonds.

Intermolecular forces

________ (IMF): weaker than either ionic or covalent bonds but are responsible for determining the state of matter a molecular compound is at a given temperature and a molecules melting point and boiling point.

Hybridization

________: several atomic orbitals mix to form the same total number of equivalent hybrid orbitals because their bond lengths are identical and their bond energies are identical.

London Dispersion

________ forces: the weakest of all molecular interactions.

Nonpolar covalent bond

________: when the bonding electrons are shared equally.

Metallic bonds

________: metal + metal, structured by delocalized electrons.

Sulfur hexafluoride

________ (SF6) can be written so that sulfur has twelve valence electrons.

Hydrogen bonds

________: when a hydrogen atom bonds to a very electronegative atom (Nitrogen, Oxygen, or Fluorine)

Network solids

________ (or network crystals: solids in which all of the atoms are covalently bonded to each other.

octet rule

The ________ can not be satisfied in molecules whose total number of valence electrons is an odd number.

ionic bonds

metals + nonmetals

Metallic bonds

metal + metal, structured by delocalized electrons

Covalent (molecular) bonds

nonmetal + nonmetals

Molecule

a neutral group of atoms joined together by covalent bonds

Diatomic Molecule

a molecule consisting of two atoms

Molecular Compound

a compound composed of molecules

Formula Unit

the base of an ionic compound NaCl is an ionic compound so the formula units are Na and Cl

Molecule

base unit of a molecular compound

Molecular Formula

a chemical formula of a molecular compound that shows how many atoms of each element a molecule contains

The Octet Rule

electron sharing usually occurs so that the atoms attain the electron configurations of noble gases (each having eight valence electrons)

Covalent Bond

joins two atoms together that are sharing a pair of electrons

Lone Pair

a pair of valence electrons that is not shared between atoms

Double covalent bond

a bond that involves two shared pairs of electrons

Triple covalent bond

a bond formed by sharing three pairs of electrons

Coordinate Covalent bond

a covalent bond in which one atom contributes both bonding electrons

Bond dissociation energy

The energy required to break the bond between two covalently bonded atoms

Nonpolar covalent bond

when the bonding electrons are shared equally

Polar covalent bond

the electrons are shared unequally

Dipole

in a polar molecule, one end of the molecule is slightly negative and the other slightly more positive

molecular orbitals

orbitals that apply to the entire molecule when two atoms combine and their orbitals overlap

Bonding orbital

a molecular orbital that can be occupied by two electrons of a covalent bond

Sigma bonds

formed when two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis and connects two atomic nuclei

Pi bonds

the second bond area formed in a double bond

hybridization

several atomic orbitals mix to form the same total number of equivalent hybrid orbitals because their bond lengths are identical and their bond energies are identical

inter

between

Intra

within

Intermolecular forces (IMF)

weaker than either ionic or covalent bonds but are responsible for determining the state of matter a molecular compound is at a given temperature and a molecules melting point and boiling point

Vander Waals Forces

attractions between the molecules that are the two weakest attractions between molecules

Dipole/Dipole interactions

when polar molecules are attracted to each other

London Dispersion forces

the weakest of all molecular interactions

Hydrogen bonds

when a hydrogen atom bonds to a very electronegative atom (Nitrogen, Oxygen, or Fluorine)

network solids (or network crystals

solids in which all of the atoms are covalently bonded to each other