AP Chemistry
Na+1
Na
Cl-1
+
+
Sodium's valence electron is
TRANSFERRED
to Chlorine!
Na + Cl
Na+ + Cl
-
+ = Sodium Ion
-
= Chloride Ion
1:1 ion ratio
NaCl
Formula Unit!
If atoms come together and bond, there
should be a net decrease in energy
the bonded state is more stable;
therefore, at a lower energy level
Formation of ions alone is not itself
energetically favorable!
IE of Na (remove an electron) = +496 kJ/mol
EA of Cl (add an electron) = -349 kJ/mol
The bond releases more than enough
energy to make it favorable and worth the
"energy investment"
Using Coulomb's Law and bond length, it is
calculated to be:
Unit 6
AP Chemistry
Additional energy is released with the
formation of the crystal lattice:
lattice energy is the change in energy that
occurs when an ionic solid is separated
into isolated ions in the gaseous phase
+786 kJ/mol for NaCl
So, to form NaCl from Na and Cl:
Fig. 9.2
pg. 332
Finding the lattice energy experimentally is difficult, so it
can be calculated using the Born-Haber Cycle, which is
based on Hess's Law.
To obtain the lattice energy of NaCl, you need to envision
2 routes that it can be formed from its elements:
Using the enthalpy tables,
Hf = -411 kJ/mol
Unit 6
AP Chemistry
IONIC COMPOUNDS
- High melting points
- Crystalline Solids at room temp.
- Metals and Nonmetals (ions)
- Formula Unit (fu!) shows ratio of ions
- Brittle due to repulsion of like
charges of ions in contact when one
layer slides across another after
being struck (shattered)
- Low vapor pressure due to strong
Coulombic interactions of ions
- When molten or dissolved in
water, conduct electricity
Electrolyte
IONIC COMPOUNDS
All ions are linked together in the crystalline
solid, held strongly by Coulombic forces of
attraction. Ions with higher charges lead to
higher Coulombic forces of attraction, and
thus, a higher melting point:
Unit 6
AP Chemistry
For ions of a given charge, the smaller the ions, the
smaller the distance between ion centers, the
stronger the Coulombic forces of attraction and thus,
a higher melting point.
Ionic Radii
Cations get smaller as they typically lose a principal
energy level; if not, as e- are lost, the Zeff increases and
pulls the e- cloud closer to the nucleus
Anions get larger as the e-/e- repulsion is increased, so
orbitals expand; or, the same Zeff is acting on more
e- now, so e- cloud not pulled as tightly inward
NaF = 993 ℃
KF = 858 ℃
RbF = 795 ℃
NaCl = 801 ℃
KCl = 770 ℃
RbCl = 718 ℃
NaBr = 747 ℃
KBr = 734 ℃
RbBr = 693 ℃
Isoelectronic
Different ions having the same number
and configuration of electrons
Unit 6
AP Chemistry
Electron Configuration of Ions
Directly related to Ionization Energy!
Group IA
Group IIA
Group IIIA
+1 lose s1 electron
+2 lose s1 & s2 electrons
+3 lose s1,s2 & p1 electrons
Al, Ga, In, Tl
+1 lose p1 electron Ga, In, Tl
Losing successive electrons requires more
energy...Group IA and IIA primarily form
ionic compounds, group IIIA forms less
Electron Configuration of Ions
Group IVA
+2 lose s1 & s2 electrons Sn, Pb
+4 lose s1,s2 & p2 electrons Sn, Pb
C, Si, Ge = no ions
Group VA
-3 gain 3e- to be ns2np6 N, P, As
+3 lose p3 electrons As, Sb, Bi
+5 lose p3 & s2 electrons As, Sb, Bi
Group VIA
-2 gain 2e- to be ns2np6 (not Po)
Group VIIA
-1 gain 1e- to be ns2np6 (not At)
Transition metals commonly form a +2 ion by losing their
s1 & s2 electrons; they can also form other ions by losing
d electrons as well: it all depends on what is energetically
favored based on electron configurations.
Inner-transition metals commonly form a +3 ion by losing
their s1, s2 and d1 electrons.
Unit 6
AP Chemistry
MOLECULAR COMPOUNDS
- Low melting points
- Mainly liquids and gases at room temp.
a few solids - ex. table sugar,
but melt at very low temperatures
- ONLY nonmetals
- Molecular Formulas give actual numbers
of atoms in the molecule (C12H22O11)
- When molten or dissolved in water, do
not conduct electricity, or do very little
Non-electrolyte / Weak-electrolyte
nonpolar
polar
Covalent Bond
the force of attraction between oppositely charged
nuclei and electrons holding atoms together
Bond length:
Bond Dissociation Energy
74 pm
Bond Length
Bond Dissociation Energy:
Unit 6
AP Chemistry
H + Cl
H Cl
Coordinate Covalent Bond
a bond formed when both electrons of
the bond come from the same atom
Unit 6
Ammonium
Hydronium
NH41+
H3O1+
AP Chemistry
Octet Rule
the tendency of atoms in molecules to
have 8e- in their valence shell (H=2)
Multiple Bonds
Single bond - single pair of e- shared
Double bond - two pairs of e- shared
Triple bond - three pairs of e- shared
So how do we know what type of bond will form between elements?
It will depend on their
ELECTRONEGATIVITIES!
Unit 6
O-H = polar
any-F = polar
C-H = nonpolar
C-O = polar
C-Cl = polar
C-S = nonpolar
N-O = polar
H-Cl = polar
C-N = nonpolar
O-S = polar
H-N = polar
AP Chemistry
Notice that the larger the difference between the electronegativities, the
stronger the bond.
Also notice the polar covalent electron cloud picture. The dipole is represented by the symbols
δ+ and δ- meaning partially positive or negative (that is the lowercase Greek letter delta).
*only seen on Polar Bonds! see picture above!
Lewis Electron-Dot Formulas
Unit 6
AP Chemistry
Lewis Electron-Dot Formulas
1. Calculate the total number of valence electrons:
- if it is a polyatomic anion, add the charge in
- if it is a polyatomic cation, subtract the charge
2. Draw the skeleton structure - the central atom is
typically the least electronegative atom - connect
the rest of the atoms with single bonds
3. Distribute the electrons to the atoms surrounding
the central atom/atoms to satisfy the octet rule
4. Distribute the remaining electrons as pairs to the
central atom/atoms - if there are fewer than 8e- on
a central atom, this suggests a double/triple bond!
COCl2
SCl2
NF4
Unit 6
1+
AP Chemistry
Exceptions to the Octet Rule
some molecules have an atom
with fewer than 8 valence eGroup IIA (Be) or IIIA (B, Al)
BeF2
VE = 24
Cl Al Cl
Cl
BF3
AlCl3
Cl
Cl
Al2Cl6
Al
Cl
Cl
Al
Cl
Cl
some molecules have an atom with
more than 8 valence e- ( fairly common)
*Period 3 and on, available d orbitals!
PF5
Unit 6
XeF4
AP Chemistry
Resonance
The bonding electrons of a covalent bond are not
always located between the two bonded atoms...
delocalized bonding: the bonding pair of electrons
is spread over a number of atoms rather than
localized between two atoms
When this occurs, we cannot use just a single
electron-dot formula, so we need to show...
RESONANCE: showing all possible electron-dot
formulas and/or combination of them
Resonance
Common Examples:
Ozone
Carbonate Ion
The double-arrow notation
does not mean that the
molecule "flips" back-andforth between the
resonance structures...there
is only one molecule!
Benzene
Unit 6
AP Chemistry
Delocalized bonding also occurs in metals:
metallic bonding! Positive metal nuclei (ions)
surrounded by a "sea of electrons" giving
metals there very important characteristics:
shiny, malleablity, ductility, conductivity
Unit 6
AP Chemistry
Formal Charge:
You can use the formal charge concept to help
determine the best Lewis formula for a molecule!
COCl2
Rules for formal charge:
1. Half of the electrons of a bond are assigned to
each atom in the bond (each dash = 2e-)
2. Both electrons of a lone pair are assigned to the
atom to which the lone pair belongs
3. Formal charge =
#valence e- - (#bonds + #lone pair e-)
Sum of formal charges = charge of molecular species
Ex: if polyatomic ion has a charge of -1,
formal charges need to add up to -1
Formal charge = #valence e- - (#bonds + #lone pair e-)
COCl2
So which structure is better?
A. The one with the lowest magnitudes of formal charges
B. If two options have the same formal charge
magnitudes, chose the one having the negative
formal charge on the more electronegative atom
C. When possible, choose the Lewis structure that does
not have like charges on adjacent atoms
Unit 6
AP Chemistry
Formal Charge:
You can use the formal charge concept to help with
writing the skeleton structure of a molecule!
VE = 26
SOCl2
So which structure is better?
A. The one with the lowest magnitudes of formal charges
B. If two options have the same formal charge
magnitudes, chose the one having the negative
formal charge on the more electronegative atom
C. When possible, choose the Lewis structure that does
not have like charges on adjacent atoms
Bond Length
the distance between the nuclei in a bond
*determined experimentally using x-ray
diffraction and/or molecular spectra
Bond lengths for covalent single bonds
can be predicted from covalent radii:
Add-up the radii of the two atoms!
Table 9.4, pg. 358
C-Cl
N-O
H-P
Unit 6
AP Chemistry
Bond Order
the number of bonds between atoms
(bonds/bond regions for molecules)
1
C C
3
C C
C C
*2
exhibiting resonance will have
* Molecules
bond orders of decimals like 1.33 or 1.5, for
example:
O O
132 pm
O
O O O
C
O C O
123 pm
114 pm
bond order =
C O
143 pm
133 pm
123 pm
bond order =
*resonance bonds will have a fractional bond order,
but a fractional bond order does NOT mean the
molecule HAS TO exhibit resonance!
Unit 6
AP Chemistry
Bond Energy
the average enthalpy change for the breaking
of a bond in a molecule in the gas phase
Bond Dissociation Energy
9.5
ble 360
a
T .
pg
In general, the
H of a reaction is equal to...
Usually obtained from thermochemical
data, as they are known more accurately!
broken bond energies
Reactants!
CH4 (g) + Cl2 (g)
formed bond energies
Products!
CH3Cl (g) + HCl (g)
[BE(C-H) + BE(Cl-Cl)] - [BE(C-Cl) + BE(H-Cl)]
Unit 6
AP Chemistry
Bond Energy
Most useful for...
Explaining heats of reaction:
In general, a reaction is EXOTHERMIC if...
In general, a reaction is ENDOTHERMIC if...
Understanding stabilities of compounds:
Complex of nitrogen triiodide
and ammonia is"contact explosive"
as it is so sensitive...
nitrogen-iodine single bonds
are replaced by very stable
nitrogen-nitrogen triple bonds
and iodine-iodine single bonds
NI3-NH3 and a feather!
Lewis Dot Structures
provide us a flat,
2-dimensional look at
molecules.
VSEPR Theory predicts
shapes of molecules &
ions based on e- pairs
arranging to minimize
e- pair repulsions
Unit 6
AP Chemistry
Arrangement of e- pairs gives us the
ELECTRON GEOMETRY
Ideal geometries: central atom surrounded by
bonds (bonding electrons, NO lone pairs)
What we "see" = nuclear positions =
Molecular Shape!
Unit 6
AP Chemistry
What we "see" = nuclear positions =
MOLECULAR SHAPE!
non-bonding electrons alter the ideal geometry
See-Saw
Based on Molecular Shape, we can determine
whether or not a MOLECULE will be
Polar or Nonpolar
Do the dipole moments cancel out or converge?
Table 10.1, pg. 387
Can be measured experimentally-evidence of the shapes!
Water = bent, not linear because it is polar!
it has a net dipole moment
Unit 6
AP Chemistry
VSEPR predicts molecular geometry
It is good, but has its limits.
To understand bonding...
Two theories we will look at:
Valence Bond Theory
(orbital hybridization)
*some controversy here! Not
everyone accepts spd
hybridization...
sp3d and sp3d2 are NOT
on the AP Exam!
Shapes, yes! Proven by NMR!
Unit 6
AP Chemistry
Molecular Orbital Theory
Superposition of atomic orbitals...a.k.a.
linear combination of atomic orbital wave functions
Unit 6
AP Chemistry
Sigma and Pi Bonds
Bond Rotation
Sigma: allows rotation
Unit 6
Pi: inhibits
AP Chemistry
Ethene
H2C=CH2
How many
Unit 6
Ethyne
HC=CH
bonds and how many
bonds?
AP Chemistry
Unit 6
AP Chemistry
Use simple structure and bonding models to
account for each of the following:
a) the bond length between 2 nitrogen
atoms is longer in N2H2 than in N2
b) all the bond lengths in CO32- are identical
and longer than a carbon-oxygen double bond
Unit 6
AP Chemistry
In the 3D molecule, angle y is not really 90o like it
appears to be in the drawing. Explain why in terms of
electron domains (VSEPR).
In the 3D molecule, angle x is not really 180o like it
appears to be in the drawing. Estimate what the angle
truly is and justify your answer.
Unit 6
AP Chemistry
ELECTRONEGATIVITIES!
Unit 6