chapter 9 covalent bonding. covalent bond sharing of electrons –nonmetal- nonmetal –...
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Covalent bond
• Sharing of electrons– Nonmetal- nonmetal– electronegativity
difference less than 1.7
2 types of covalent bonds
• Polar covalent- bonding electrons are shared unequally
• Nonpolar covalent- bonding electrons are shared equally– Balanced distribution of charge
How do you tell the difference in polar and nonpolar?
• If the 2 atoms have a difference in electronegativity the bond is polar, if no difference then the bond is nonpolar
• If polar bonds draw the electron to one side of the molecule and the other side has none the molecule is polar
Bond length
• Distance between two bonded atoms at their minimum potential energy– Average distance between bonded atoms
Nomenclature for covalent bonds
• (nonmetal- nonmetal)
• 2 systems: – Stock- name(+), (+) oxidation # in roman
numerals, name(-)– Classical- use prefixes except for a single (+)
ion • *Never use mono first*
Lewis structures
• Use electron dot diagrams to show bonding and electron arrangement
Symbols uses in molecular structural formulas
unshared pair- (lone pair) pair of electrons that is not involved in bonding- belongs to one atoms
single bond- 2e- shared
double bond- 4e- shared
triple bond- 6e- shared
Structural formula
• Indicates kind, number, arrangement, and bond type in a molecule
VSEPR theory
• Valance shell electron pair repulsion– Model for molecular geometry– Bond angles– Arrangement minimizes repulsion of e- around
the central atom– Molecules adjust their shape, so that valence
e- are as far apart as possible
Hybrid orbitals
• Orbital of equal energy produced by the combination of two or more atomic orbitals
• (sp, sp2, sp3)
Dipole moment
• Measure of the strength of the dipole and is a property that results from the asymmetrical charge distribution in a polar molecule- depends on size and distance Qd
Hydrogen bonding
• Intermoleculer (Van der Waals) force in which H bonded to a highly electronegative atom is attracted to an unshared pair of an electronegative atom in a nearby molecule
London Dispersion forces
• (dispersion force) result from the constant motion of e- ’s and the creation of instanteous dipoles– Force generated in a temporary dipole
interaction– Most important Van der Waals force– Proposed by Fritz London in 1930– Strength increases with the number of e- in
the interacting atoms
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