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31 Cards in this Set

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Ionic bond
complete transfer of e- from atom w/ smaller ionization energy to atom w/ greater ionization energy affinity, ions held together by electrostatic forces.

difference of electronegativity greater than 1.7

Group I/II + Group VII

high MP
high BP
conduct electricity in liquid & aq. states, not solid

ionic solids form from crystal lattices
Covalent bonds
sharing e- btw 2 or atoms with similar electronegativity.

energy req. to form ions is greater than the energy that would be released upon the formation of an ionic bond.

low MP
do not conduct electricity in liq or aq states.
Bond order
# of shared e- pairs btw 2 atoms
1, 2, 3 e- pairs = single, double, or triple covalent bond.
Covalent bond characterization
1. Bond length: avg length btw 2 nuclei of atoms involved in bond
triple bond < double bond < single bond.

2. Bond energy: energy req to separate 2 bonded atoms

str of bond increase as # of shared e- pair increases.
triple bond > double bond > single bond.
Bonding e-
shared valence e- of covalent bond
Non-bonding e-
valence e- not involved in covalent bond

unshared e- pairs = lone e- pairs
Lewis structures
-Lewis dot symbol
-least electronegative atom is central atom.
-count all valence e- of atoms
formal charge
V = 1/2 Bonding electrons - non bonding electrons
Resonance
A Lewis structure w/ small or no formal charge is preferred over a Lewis Structure w/ large large formal charge.

Lewis structure /w neg. formal charge place on more electronegative atom, more stable
Exceptions of octet rule
Atoms found in or beyond 3rd period can have more than 8 valence e-, occupy d orbitals
Polar covalent bond
occurs btw atoms with small differences 0.4 to 1.7 Pauling units

More electronegative atom is partially Negative

Less electronegative atom is partially positive
Dipole Moment
vector quantity cursive-m, defined as the product of the charge magnitude (q) and the distance btw 2 partial charges (r): cursive m = q*r
-denoted by arrow pointing from pos to neg charge measured by Debye units (coulomb-meters)
Non-covalent bond
occurs btw atoms that have same electronegativity

nonpolar covalent bonds occur diatomic molecules H2, Cl2, O2, N2
Coordinate covalent bond
share e- pair, from lone pair of one of atoms in molecule, typically found in Lewis acid-base compounds
Lewis (a)cid
compound that can (a)ccept e- pair to form covalent bond

F H
| |
ex: F--B + :N--H
| |
F H
Lewis acid Lewis base
Lewis base
compound that can donate an e- pair to form covalent bond
Valence Shell Electron-Pair Repulsion Theory (VSEPR theory)
3-D arrangement of atoms surrounding central atom
-determine by repulsion btw bonding & nonbonding e- pairs
-e- pairs arrange themselves far apart as possible minimizing repulsion
Linear
180 degrees
Trigonal Planar
AX_3 = 120 degrees
Tetrahedral
AX_4 = 109.5 degrees
Trigonal Pyramidal
AX_3E_1 < 109.5

..
|
A
/ | \
X X X
Trigonal Bipyramidal
AX_5 = 90, 120, 180 degrees

X
x |
`A- x
/ |
x X
See Saw
AX_4E > 90 degrees
< 120 degrees

..
x |
`A-x
/ |
x x
Octahedral
AX_6
Bent
AX_2E_1 < 120 degrees

AX_2E_2 < 109.5 degrees
Polarity of Molecules
-molecule w/ net dipole moment is polar
-molecule w/ nonpolar bonds is always nonpolar
-molecule w/ polar bonds may be polar or nonpolar depending on orientation of bonding dipoles
ex:

Cl
|
C
/ | \
Cl Cl Cl
Molecular orbital
+ + = bonding orbital

+ - = antibonding orbital

sigma bonds = orbitals overlap head to head

pi bonds = orbitals are parallel
Van der Waal forces
dipole-dipole interactions
hydrogen bonding
dispersion forces
Dipole-Dipole interaction
-orient (+) region of 1 molecule close to (-) region of another molecule
-attractive forces only effective close
-dipole-dipole present in solid & liquid, not gas
-polar increase BP
Hydrogen bond
-Strong dipole-dipole
-can be either INTRA or INTERmolecular
-Increases BP, more energy to break H-Bond
-important in behavior of water, alcohols, amines, & carboxylic acids
Dispersion forces
-Present in all molecules

-bonding e- of covalent bond may be equally shared, but will be located randomly throughout the orbit
-Permit unequal sharing of e- causing rapid polarization and counter polarization of e- cloud & short lived dipole = London forces

-London forces = temporary attractive force from e- from 2 adjacent atoms occupy position that make the atoms form temporary dipoles a.k.a. induced dipole.
-generally weaker than other intermolecular forces
-works in close proximity btw atoms
-large molecules, e- are far from nucleus are relatively easy to polarize therefore posses greater dispersion forces