Chem II: 1-2 Flashcards
octet rule
atom tends to bond with other atoms so that it has 8 electrons in outermost shell –> forming a stable electron configuration similar to that of noble gases
exceptions to octet rule
incomplete octet
stable with fewer than 8 electrons
H (2), He (2), Li (2), Be (4), B (6)
exceptions to octet rule
expanded octet
any elemet in period 3 and greater can hold more than 8 electrons
exceptions to octet rule
odd numbers of electrons
any molecule with odd number of valence electrons cannot distribute those electrons to give 8 to each atom
ex: NO
atoms that almost always abide by octet rule
C, N, O, F, Na, Mg
nonmetals ____ electrons
gain
metals ____ electrons
lose
ionic bonding
one or more electrons from an atom with low ionization energy (typically metals) are transferred to atom with high electron affinity (typically nonmetals)
resulting in electrostatic attraction between opposite chages
coordinate covalent
if both of shared electrons in covalent bond are contributed by only one of the 2 atoms
difference between ionic and covalent compounds
ionic - gain/loss of electrons
covalent - sharing of electrons
alkali and alkaline earth metals readily form ionic bonds with…
halogens of group VIIA
physical characteristics of ionic compounds
- high MP and BP due to electrostatic attractions
- solubility of ions in water due to interactions with polar solvents
- good conductors of heat and electricity
- crystal lattice arrangement to minimize repulsive forces
- large electronegativity differences between ions
why do ionic bonds tend to form between metals and nonmetals
metals lose electrons because they have low ionization energies
nonmetals gain electrons because they have high electron affiinities
physical characteristics of covalent compounds
- relatively week intermolecular interactions
- lower MP and BP
- poor conductor of electricity bc do not break down into constituent ions
bond length
avg distance between the two nuclei of atoms in a bond
as the number of shared electron pairs increases, bond length ______ bc…
decreases because the two atoms are pulled closer together
bond energy
energy required to break a bond by separating its components into their isolated, gaseous atomic states
the greater the number of pairs of electrons shared between the atomic nuclei, the _____ energy is required to break the bonds holding the atoms together
more
7 common diatomic molecules are:
(make a 7 in periodic table)
H2, N2, O2, F2, Cl2, Br2
dipole moment eq
p = qd
p: dipole moment (debye units –> coulomb meters)
q: magnitude of charge
d: distance
what kinds of rxns are coordinate covalent bonds usually found?
nucleophile-electrophile rxns, lewis acid-base rxns, complexation rxns
nonbonding electrons
electrons in valence shell that are not involved in covalent bonds
bonding electrons
valence electrons involved in a covalent bond
formal charge
assumes equal sharing of all bonded electron pairs
if the possible lewis structures differ in their bond connectivity or arrangement, then the lewis structures represent ______
different possible compounds
if the lewis structures show the same bond connectivity and differ only in the arrangement of the electron pairs, then these structures represent _____
different resonance forms of a single compounds
formal charge - most stable compounds
minimizes the number and magnitude of formal chages
steps in drawing lewis dot structure
- central atom - least electronegative
- hydrogen and halogens (F, Cl, Br, I) usually at ends
- usually
- C - 4 bonds
- O - double bond
- F - 1 bond
how to calculate formal charge
formal charge = valence electrons - nonbonding electrons - 1/2 bonding electrons
formal charge ____estimates the effect of electronegativity differences, while ox numbers ___estimate the effect of electronegativity differences
under
over
resonance
allows for greater stability, delocalizing electrons and charges over pi system
the more stable the structure, the ____ it contributes to the character of the resonance hybrid
more
using formal charges to assess the stability of resonance structures
lewis structures
- small or no formal charges is preferred over more
- less separation between opposite charges is preferred over larger
- negative formal charges are placed on more electronegative atoms is more stable
linear
2 things, no lone pair
6 things, 4 lone pairs
trigonal planar
3 things, no lone pairs
tetrahedral
4 things, no lone pairs
trigonal bipyramid
5 things, no lone pairs
octahedral
6 things, no lone pair
trigonal pyramid
4
1 lone pair
seesaw
5
1 lone pair
square pyramid
6
1 lone pair
VSEPR?
2
linear
VSEPR?
3
trigoonal planar
VSEPR?
4
tetrahedral
VSEPR?
5
trigonal bipyramid
VSEPR?
6
octahedral
VSEPR?
3
1 lone pair
less than 120
VSEPR?
4
1 lone pair
VSEPR?
5
1 lone pair
seesaw
VSEPR?
6
1 lone pair
4 things
2 lone pairs
<<109
bent
5 things
2 lone pairs
t shaped
5 things
3 lone pairs
linear
6 things
3 lone pairs
t shaped
6 things
4 lone pairs
linear
6 things
2 lone pairs
hybridization
3 things
120 deg
3sp2
1 p leftover
hybridization
4 things
109 deg
4sp3
nothing left over
hybridization
2 things
180 deg
2sp
2 p leftover
electronic geometry
describes the spatial arrangement of all pairs of electrons around the central atom, including both the bonding and lone pairs
molecular geometry
describes the spatial arrangement of only the bonding pairs of electrons
coordination number
coordination number
number of atoms that surround and are bonded to a central atom
important for determining molecular geometry
molecular orbital
probability of finding the bonding electron sin a given space
found by combining the wave functions of the atomic orbitals
bonding orbital forms if
signs of the 2 atomic orbitals are the same
antibonding orbital forms if
the signs of the two atomic orbitals are different
sigma bond
when orbitals overlap head ot head
free rotation about axis
pi bond
two parallel electron cloud densities
no free rotation bc cannot be twisted in a way that allows continuous overlapping of the clouds of electron densities
In a hypothetical molecule, a Nitrogen has two sets of lone pair electrons and two covalent bonds. What would its formal charge be?
(A) -2
(B) -1
(C) 0
(D) +1
(B) -1
Formal Charge = 5 - [4 + 4/2] = -1
Draw the dot structure for O3, being sure to assign formal charges as needed.
Draw the dot structure for Nitrate, being sure to assign formal charges as needed.
draw the three resonance structures for Nitrate.
CRB In order to avoid having to draw multiple resonance structures, a resonance hybrid can be drawn. Which of the following are correct descriptions of a resonance hybrid?
I. Resonance hybrids often depict partial charges.
II. For where a double or single bond may exist, a dotted line can be used to represent the second bond.
III. Resonance hybrids are considered the “averages” of the possible resonance structures.
(A) I and II only
(B) I and III only
(C) II and III only
(D) I, II and III
(D) I, II and III
Each of the following are a correct description of resonance hybrids:
I. Resonance hybrids often depict partial charges.
II. For where a double or single bond may exist, a dotted line can be used to represent the second bond.
III. Resonance hybrids are considered the “averages” of the possible resonance structures.
What is the electron geometry of Sulfur Dioxide? Molecular geometry?
Electron geometry of Sulfur Dioxide is trigonal planar.
Molecular geometry of Sulfur Dioxide is bent/angular.
CRB For a tetrahedral compound, which of the following is the ideal bond angle (in degrees)?
(A) 90
(B) 104.5
(C) 109.5
(D) 120
(C) 109.5
For a tetrahedral compound, the ideal bond angle is 109.5 degrees.
What is the electron geometry of NH3? Molecular geometry?
Electron geometry of NH3 is tetrahedral.
Molecular geometry of NH3 is trigonal pyramidal.
What is the electron geometry of H2O? Molecular geometry?
Electron geometry of H2O is tetrahedral.
Molecular geometry of H2O is bent.
What is the Electron Geometry of ICl2+? Molecular Geometry?
Electron Geometry of ICl2+ is Tetrahedral.
Molecular Geometry of ICl2+ is Bent.
a) aluminum
b) calcium
c) vanadium
d) scandium
c) vanadium
a) being in very polar solvents
b) having component atoms that have an odd number of valence electrons in their elemental forms
c) being in polar aprotic solvents
d) having an odd number of valence electrons across the entire molecule
d) having an odd number of valence electrons across the entire molecule
c
b) PCl3
a
a
c) 2
b) 0
c) 16
intermolecular forces from weakest to strongest
london disperson forces/van der waals < dipole-dipole interactions < hydrogen bonding
london dispersion forces
- type of van der waals force
- rapid polarization and counterpolarization of electron cloud and formation of short-lived dipole moments
- weakest of intermolecular interactions
- do not extend over long distances
- relevant only when molecules are in close proximity
- strength depends on degree and ease by which molecules can be polarized
- large molecules more easily polarizable
dipole dipole interactions
- present in solid and liquid phases
- negligible in gas phase bc of inc distance between gas particles
- why polar species have higher MP and BP
hydrogen bonding
- strongest
- when H is bonded to NOF (very electronegative)
- may be intra or inter moelcualr
- no actual sharing or transferring of electrons between two atoms
- unusually high BP and MP
What is the Electron Geometry of SF4? Molecular Geometry?
Electron Geometry of SF4 is Trigonal Bipyramidal.
Molecular Geometry of SF4 is See-saw.
What are the bond angles for Trigonal Bipyramidal? See-saw?
T-structure? Linear?
What are the bond angles for Octahedral? Square Pyramidal? Square Planar?
a) ClF3
b) H2O
c) CH4
d) XeF2
d) XeF2
