Download Chemistry 1 Lectures

Brad Collins, DTCC
Chemical Structures
(Part 2)
Chapter 11
Some images Copyright © The McGraw-Hill Companies, Inc.
Organic Structures
•
Organic compounds do not follow the AXn general
formula, but:
•
DO generally follow the octet rule.
•
Hydrogen and halogens are always outer atoms
•
In neutral atoms, N and P will retain their lone
pair; O and S will retain their 2 lone pairs
•
Rarely see expanded octets
11.1
Organic Compounds
•
Use structural formulas to depict organic
compounds
•
Isomer - compounds with the same molecular
formula, but different structural formulas.
•
Ethanol:
•
Dimethyl ether:
•
Molecular: C2H6O
•
Molecular: C2H6O
•
Structural:
•
Structural:
•
Condensed structural: CH3CH2OH
•
Condensed structural: CH3OCH3
11.1
Organic Compounds
•
Condensed structural formulas
•
Hydrogens follow the atom they are bonded to.
•
•
Single bonds are omitted
•
•
CH3CH3OH, not CH3–CH3–O–H
Multiple bonds ARE shown
•
•
CH3, not H3C
CH2=CH2
or
CH
CH
Side chains are often shown:
•
but may be written in parenthesis: CH3CH(OH)CH3
11.1
Organic Compounds
•
Condensed structural formulas
•
Cyclic structures a drawn with the geometric figure they
resemble:
•
Cyclohexane, C6H12
11.1
Practice
Draw the Lewis structure for diethyl ether, CH3CH2OCH2CH3
11.1
Bond Enthalpy
•
The enthalpy required to break a particular bond in
1 mole of gaseous molecules
Bond Enthalpy
ΔH0 = 436.4 kJ
H2 (g)
H (g) + H (g)
Cl2 (g)
Cl (g) + Cl (g) ΔH0 = 242.7 kJ
HCl (g)
H (g) + Cl (g) ΔH0 = 431.9 kJ
O2 (g)
O (g) + O (g) ΔH0 = 498.7 kJ
O
O
N2 (g)
N (g) + N (g) ΔH0 = 941.4 kJ
N
N
Bond Energies
Single bond < Double bond < Triple bond
11.2
Bond Enthalpy
•
In polyatomic molecules (e.g., H2O) the average
bond enthalpy is listed
H2O (g)
H (g) + OH (g) ΔH0 = 502 kJ
OH (g)
H (g) + O (g) ΔH0 = 427 kJ
502 + 427
= 464 kJ
Average OH bond energy =
2
11.2
Bond
Enthalpy
Note: Forming the listed
bond, releases the tabled
amount of energy
•
•
Breaking bonds is
endothermic
Forming bonds is
exothermic
11.2
Bond Enthalpy (BE) and Enthalpy changes in reactions
Imagine a reaction proceeding by breaking all bonds in the
reactants and then using the gaseous atoms to form all the
bonds in the products.
ΔH0 = total energy input – total energy released
= ΣBE(reactants) – ΣBE(products)
11.2
Bond Enthalpy
and ∆Hrxn
H2(g) + F2 (g) —> 2HF(g)
!
∆Hrxn = –543.2 kJ
11.2
Practice
Use bond energies to calculate the enthalpy change for:
H2 (g) + F2 (g)
2HF (g)
Type of
bonds broken
Number of
bonds broken
Bond energy
(kJ/mol)
Energy
change (kJ)
H
H
1
436.4
436.4
F
F
1
156.9
156.9
Type of
bonds formed
H
F
Number of
bonds formed
Bond energy
(kJ/mol)
2
568.2
Energy
change (kJ)
1136.4
ΔH0 = 436.4 + 156.9 – 2 x 568.2 = -543.1 kJ
11.2
Shapes of Molecules
•
VSEPR Theory
•
Valence Shell Electron Pair Repulsion
•
Molecules will form shapes that minimize the
repulsive force between the electron pairs
(bonding pairs and lone pairs)
•
Arrangement describes how the electron pairs
arrange themselves around an atom
•
Shape describes how the molecule actually
looks in 3-dimensional space
11.3
Atomic Arrangements
Two Items Around a Central Atom
•
Classify molecules according to how many atoms (B) are bonded to
the central atom and how many lone pairs (E) are on the central atom.
•
Two outer atoms, no lone pairs (AB2)
•
Electron pairs align opposite the central atom, A
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB2
2
0
Arrangement of
electron pairs
Molecular
Shape
linear
B
linear
B
B
B
11.3
Beryllium Chloride
Cl
Be
Cl
lone pairs
on to
central
atom
2 0atoms
bonded
central
atom
10.1
Atomic Arrangements
Three Items Around a Central Atom
•
Three outer atoms, no lone pairs (AB3)
•
Class
AB3
Electron pairs align at 0, 120, and 360º around
the central atom, A
# of atoms
bonded to
central atom
3
# lone
pairs on
central atom
0
Arrangement of
electron pairs
Molecular
Shape
trigonal
planar
trigonal
planar
120º
11.3
Boron Trifluoride
Trigonal Planar
10.1
Atomic Arrangements
Three Items Around a Central Atom
•
2 outer atoms, one lone pair (AB2E)
•
Class
AB2E
Electron pairs align at 0, 120, and 360º around
the central atom, A
# of atoms
bonded to
central atom
2
# lone
pairs on
central atom
Arrangement of
electron pairs
trigonal
planar
1
B
Molecular
Shape
bent
B
11.3
Atomic Arrangements
Four Items Around a Central Atom
•
Four outer atoms, no lone pairs (AB4)
•
Electron pairs align at ~109.5º angles around the
central atom, A
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB4
4
0
Arrangement of
electron pairs
Molecular
Shape
tetrahedral
B
tetrahedral
B
B
B
11.3
Methane
11.3
Atomic Arrangements
Four Items Around a Central Atom
•
Three outer atoms, one lone pair (AB3E)
•
Electron pairs align at ~109.5º angles around the
central atom, A
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB3E
3
1
Arrangement of
electron pairs
tetrahedral
Molecular
Shape
trigonal
pyramidal
11.3
Atomic Arrangements
Four Items Around a Central Atom
•
Two outer atoms, two lone pairs (AB2E2)
•
Electron pairs align at ~109.5º angles around the
central atom, A
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB2E2
2
2
Arrangement of
electron pairs
tetrahedral
Molecular
Shape
bent
11.3
Effect of Lone Pairs
bonding-pair vs.
<
bonding
pair repulsion
lone-pair vs.
bonding
pair repulsion
<
lone-pair vs.
lone pair
repulsion
11.3
Atomic Arrangements
Five Items Around a Central Atom
•
Five outer atoms, no lone pairs (AB5)
•
Electron pairs align at ~109.5º angles around the
central atom, A
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB5
5
0
Arrangement of
electron pairs
Molecular
Shape
Trigonal
bipyramidal
Trigonal
bipyramidal
B
B
B
B
B
11.3
Phosphorus Pentachloride
11.3
Atomic Arrangements
Six Items Around a Central Atom
•
Six outer atoms, no lone pairs (AB6)
•
Electron pairs align at ~109.5º angles around the
central atom, A
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
Arrangement of
electron pairs
Molecular
Shape
AB6
6
0
Octahedral
Octahedral
B
B
B
B
B
B
11.3
11.3
Sulfur Hexafluoride
11.3
Atomic
Arrangement
Summary
11.3
Polarity
•
Definition: Unequal sharing of electrons in covalent
bonds
•
Electrons spend more time with one atom in the bond
•
Bond has negative and positive end
electron poor
region
electron rich
region
H
F
δ+
δ−
11.4
Polar Molecules
•
To be polar, molecules must meet 2 criteria:
•
Must have a polar bond (dipole)
•
Must have a molecular shape that allows a dipole moment to exist.
•
•
Dipoles are vectors (have magnitude and direction)
•
Dipole moment is the vector sum of all the dipoles present.
Examples: HF, NH3, and NF3
HF
electron poor
region
electron rich
region
H
F
δ+
δ−
11.4
NH3
NF3
11.4
Which of the following molecules have a dipole moment?
H2O, CO2, SO2, and CH4
O
H
H
dipole moment
polar molecule
S
O
O
dipole moment
polar molecule
H
O
C
O
no dipole moment
nonpolar molecule
H
C
H
H
no dipole moment
nonpolar molecule
11.4
Does CH2Cl2 have
a dipole moment?
11.4
Dipole Moments of Polar
Molecules
11.4
Trends in Polarity
•
Molecules with a linear, trigonal planar, or
tetrahedral arrangement are predicted to be
nonpolar if:
•
There are no lone pairs on the central atom
•
All the outer atoms are the same
•
If one or more of these aren’t true, the molecule is
usually polar.
•
Does NOT work for trigonal bipyramidal or
octahedral arrangements.
11.4
Practice: Polar Molecules
Predict the polarity of the following molecules:
Chloroform, CHCl3
Phosphorus trichloride, PCl3
Bromine pentafluoride, BrF5
11.4