bonds forces that hold groups of atoms together and make them function as a unit
Post on 17-Jan-2016
219 Views
Preview:
TRANSCRIPT
Bonds
Forces that hold groups of atoms together and make them function as a unit.
Bond Energy
- It is the energy required to break a bond.
- It gives us information about the strength of a bonding interaction.
Bond Length
The distance where the system energy is a minimum.
08_130
Sufficiently far apartto have no interaction
En
erg
y (k
J/m
ol)
0 Internuclear distance (nm)0.074
-458
0
(H H bond length)
HH
H H
H H
H H
(a) (b)
+
H atom H atom
The atoms begin to interactas they move closer together.
+
H atom H atom
H2molecule
+
Optimum distance to achievelowest overall energy of system
+
+
+
Ionic Bonds
- Formed from electrostatic attractions of closely packed, oppositely charged ions.
- Formed when an atom that easily loses electrons reacts with one that has a high electron affinity.
Ionic Bonds
Q1 and Q2 = numerical ion charges
r = distance between ion centers (in nm)
E = 2.31 10 J nm (19 QQ r1 2 / )
Electronegativity
The ability of an atom in a molecule to attract shared electrons to itself.
= (H X)actual (H X)expected
08_132
H2.1
Li1.0
Be1.5
Na0.9
Mg1.2
K0.8
Ca1.0
Rb0.8
Sr1.0
Cs0.7
Ba0.9
Fr0.7
Ra0.9
Sc1.3
Y1.2
La-Lu1.0-1.2
Ac1.1
Ti1.5
Zr1.4
Hf1.3
Th1.3
V1.6
Nb1.6
Ta1.5
Pa1.4
Cr1.6
Mo1.8
W1.7
U1.4
Mn1.5
Tc1.9
Re1.9
Np-No1.4-1.3
Fe1.8
Ru2.2
Os2.2
Co1.9
Rh2.2
Ir2.2
Ni1.9
Pd2.2
Pt2.2
Cu1.9
Ag1.9
Au2.4
Zn1.6
Cd1.7
Hg1.9
Ga1.6
In1.7
Tl1.8
Al1.5
B2.0
Ge1.8
Sn1.8
Pb1.9
Si1.8
C2.5
As2.0
Sb1.9
Bi1.9
P2.1
N3.0
Se2.4
Te2.1
Po2.0
S2.5
O3.5
Br2.8
I2.5
At2.2
Cl3.0
F4.0
H2.1
Li1.0
Be1.5
Na0.9
Mg1.2
K0.8
Ca1.0
Rb0.8
Sr1.0
Cs0.7
Ba0.9
Fr0.7
Ra0.9
Sc1.3
Y1.2
La-Lu1.0-1.2
Ac1.1
Ti1.5
Zr1.4
Hf1.3
Th1.3
V1.6
Nb1.6
Ta1.5
Pa1.4
Cr1.6
Mo1.8
W1.7
U1.4
Mn1.5
Tc1.9
Re1.9
Np-No1.4-1.3
Fe1.8
Ru2.2
Os2.2
Co1.9
Rh2.2
Ir2.2
Ni1.9
Pd2.2
Pt2.2
Cu1.9
Ag1.9
Au2.4
Zn1.6
Cd1.7
Hg1.9
Ga1.6
In1.7
Tl1.8
Al1.5
B2.0
Ge1.8
Sn1.8
Pb1.9
Si1.8
C2.5
As2.0
Sb1.9
Bi1.9
P2.1
N3.0
Se2.4
Te2.1
Po2.0
S2.5
O3.5
Br2.8
I2.5
At2.2
Cl3.0
F4.0
Increasing electronegativity
De
crea
sing
ele
ctro
neg
ativ
ity
Increasing electronegativity
De
crea
sing
ele
ctro
neg
ativ
ity
(a)
(b)
Polarity
A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.
+
FH
08_131
F
H
F
H
F
HF
H
F
H
(a)
H F
(b)
H F
H F
H F
H F
08_133
H
O
H
(a)
+
(b)
08_134
HH
N
H
3
(a)
+
(b)
Achieving Noble Gas Electron Configurations (NGEC)
Two nonmetals react: They share electrons to achieve NGEC.
A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC.
Isoelectronic Ions
Ions containing the the same number of electrons
(O2, F, Na+, Mg2+, Al3+)
O2> F > Na+ > Mg2+ > Al3+
largest smallest
Lattice Energy
The change in energy when separated gaseous ions are packed together to form an ionic solid.
M+(g) + X(g) MX(s)
Lattice energy is negative (exothermic) from the point of view of the system.
Formation of an Ionic Solid
1. Sublimation of the solid metal
M(s) M(g) [endothermic]
2. Ionization of the metal atoms
M(g) M+(g) + e [endothermic]
3. Dissociation of the nonmetal 1/2X2(g) X(g) [endothermic]
Formation of an Ionic Solid(continued)
4. Formation of X ions in the gas phase:
X(g) + e X(g) [exothermic]
5. Formation of the solid MX
M+(g) + X(g) MX(s) [quite
exothermic]
Lattice Energy = k( / )QQ r1 2
Q1, Q2 = charges on the ions
r = shortest distance between centers of the cations and anions
Models
Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.
Fundamental Properties of Models
- A model does not equal reality.
- Models are oversimplifications, and are therefore often wrong.
- Models become more complicated as they age.
- We must understand the underlying assumptions in a model so that we don’t misuse it.
Bond Energies
Bond breaking requires energy (endothermic).
Bond formation releases energy (exothermic).
H = D(bonds broken) D(bonds formed)
energy required energy released
Localized Electron Model
A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.
Localized Electron Model
1. Description of valence electron arrangement (Lewis structure).
2. Prediction of geometry (VSEPR model).
3. Description of atomic orbital types used to share electrons or hold long pairs.
Lewis Structure
- Shows how valence electrons are arranged among atoms in a molecule.
- Reflects central idea that stability of a compound relates to noble gas electron configuration.
Comments About the Octet Rule
- 2nd row elements C, N, O, F observe the octet rule.
- 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.
- 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.
- When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.
Resonance
Occurs when more than one valid Lewis structure can be written for a particular molecule.
These are resonance structures. The actual structure is an average of the resonance structures.
Formal Charge
The difference between the number of valence electrons (VE) on the free atom and the number assigned to the atom in the molecule.
We need:
1. # VE on free neutral atom
2. # VE “belonging” to the atom in themolecule
Formal Charge
Not as good Better
CO O(-1) (0) (+1)
CO O(0) (0) (0)
VSEPR Model
The structure around a given atom is determined principally by minimizing electron pair repulsions.
Predicting a VSEPR Structure
1. Draw Lewis structure.
2. Put pairs as far apart as possible.
3. Determine positions of atoms from the way electron pairs are shared.
4. Determine the name of molecular structure from positions of the atoms.
top related