Chapter 1
Remembering
General Chemistry:
Electronic Structure
and Bonding
Paula Yurkanis Bruice
University of California,
Santa Barbara
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What is Organic Chemistry?
• Organic compounds: from living organisms
(with a vital force)
• Inorganic compounds: from minerals
(without a vital force)
organic chemistry = compounds that contain carbon
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What Makes Carbon So Special?
• Atoms to the left of carbon give up electrons.
• Atoms to the right of carbon accept electrons.
• Carbon shares electrons.
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The Structure of an Atom
Protons are positively charged.
Neutrons have no charge.
Electrons are negatively charged.
Atomic number = # of protons
Atomic number of carbon = 6
Neutral carbon has six protons
and six electrons.
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Atomic Symbol Convention
FORMULA:
1. Z = number of protons (atomic number)
2. A = number of protons + number of neutrons (mass number)
A = Z + number of neutrons
(since Z= number of protons or atomic number)
3. number of neutrons = A – Z
4. number of protons = number of electrons
Note: a. applicable for neutral atoms only
b. number of neutrons will change
5. charge of an ion = number of protons – number of electrons
Note: a. applicable for monoatomic ions only
b. number of protons is not equal to the number of electrons
c. number of electrons will change
NOTE: The number of protons will not change in a reaction but only the
number of electrons
Isotopes
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Learning Check
Naturally occurring carbon consists of three isotopes, C-12, C-13,
and C-14. State the number of protons, neutrons, and electrons
in each of the following.
12
6
protons
neutrons
electrons
C
______
______
______
13
6
______
______
______
C
14
6
______
______
______
C
Solution
12
6
C
protons
6 p+
neutrons 6 n
electrons 6 e–
13
6
C
6 p+
7n
6 e–
14
6
C
6 p+
8n
6 e–
Do You Understand Ions?
How many protons and electrons are in
27 3+
13 Al
?
13 protons, 10 (13 – 3) electrons
How many protons and electrons are in
78
2Se
?
34
34 protons, 36 (34 + 2) electrons
PROBLEMS
Answers
The Distribution of Electrons in an Atom
• The first shell is closest to the nucleus.
• The closer the atomic orbital is to the nucleus,
the lower its energy.
• Within a shell, s < p.
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ELECTRON CONFIGURATION
• Mnemonics of electron distribution
• Describes the arrangement of electrons by sublevel
according to increasing energy
Order of orbitals
(filling) in multielectron atom:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s
CONVENTION WHEN WRITING ELECTRON
CONFIGURATION
Aufbau principle: An electron goes into the atomic orbital with the lowest energy.
Pauli exclusion principle: No more than two electrons can be in an atomic orbital;
the two electrons must be of opposite spin
Hund’s rule: An electron goes into an empty degenerate orbital rather than pairing up.
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Relative Energies of the Atomic Orbitals
Following Aufbau Principle:
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Condensed Electron Configurations
Accounts for the outermost or valence electrons of an atom
• The electron configuration can be abbreviated by
indicating the innermost electrons with the symbol of
the corresponding noble gas
Note:
1. d-block = period # - 1
2. f-block = period # - 2
Core electrons are electrons in inner shells.
Valence electrons are electrons in the
outermost shells.
The chemical behavior of an element depends
on its electronic configuration.
Lewis Symbols
G. N. Lewis developed a method to denote potential
bonding electrons by using one dot for every valence
electron around the element symbol.
When forming compounds, atoms tend to gain, lose, or
share electrons until they are surrounded by eight
valence electrons (the octet rule).
Valence Electrons
Number of valence electrons = Group number
Problems
Answers
Answers
Atoms on the Left Side of the
Periodic Table Lose an Electron
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Atoms on the Right Side of the
Periodic Table Gain an Electron
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A Hydrogen Atom Can Lose or
Gain an Electron
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The Three Models of Chemical Bonding
An Ionic Bond is Formed by the Attraction
Between Ions of Opposite Charge
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Covalent Bonds are Formed by
Sharing Electrons
Nonpolar covalent bond = bonded atoms are the same
Polar covalent bond = bonded atoms are different
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How Many Bonds Does
an Atom Form?
Lewis Structures
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Bond Polarity Depends on
the Difference in Electronegativity
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Bond Polarity and Electronegativity
Difference in electronegativity is a gauge of bond
polarity:
1. Polar Covalent Bond
If the electronegativity difference between two bonded
atoms is > 0 - 1.9 (unequal sharing of electrons)
2. Ionic Bond
If the electronegativity difference is 2.0 or more (transfer
of electrons).
3. Nonpolar Covalent Bond
If the electronegativity difference is 0 (equal sharing of
electrons);
Bond Polarity and Electronegativity
Example:
Which would be more polar, a H-F bond or a H-Cl
bond?
H-F 4.0 - 2.1 = 1.9
H-Cl 3.0 - 2.1 = 0.9
Which of the following is polar?
1. N – N
2. S – Cl
3. Na – F
Classify the following bonds as ionic, polar covalent,
or nonpolar covalent: The bond in CsCl; the bond in
H2S; and the NN bond in H2NNH2.
Cs – 0.7
Cl – 3.0
3.0 – 0.7 = 2.3
Ionic
H – 2.1
S – 2.5
2.5 – 2.1 = 0.4
Polar Covalent
N – 3.0
N – 3.0
3.0 – 3.0 = 0
Nonpolar Covalent
Polar Covalent Bonds
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Problems
Answers
Dipole Moments
the greater the difference in electronegativity,
the greater the dipole moment, and the more polar the bond
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Electrostatic Potential Maps
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Formal Charge
Formal Charge = the number of valence electrons
– (number of lone pair electrons +1/2
number of bonding electrons)
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Neutral Oxygen Forms Two Bonds
Oxygen has two lone pairs.
If oxygen does not form two bonds, it is charged.
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Neutral Nitrogen Forms Three Bonds
Nitrogen has one lone pair.
If nitrogen does not form three bonds, it is charged.
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Neutral Carbon Forms Four Bonds
if carbon does not form four bonds, it has a charge
(or it is a radical)
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Hydrogen and the Halogens
Form One Bond
A halogen has three lone pairs.
if hydrogen or halogen does not form one bond, it has a charge
(or it is a radical)
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The Number of Bonds and Lone Pairs
• Halogen = 3 lone pairs
• Oxygen = 2 lone pairs
• Nitrogen = 1 lone pair
In order to have a complete octet, the number of bonds and the number of lone pairs
must total four.
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Lewis Structures
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How to Draw a Lewis Structure
NO3–
Determine the total number of valence electrons (5 + 6 + 6 + 6 = 23).
Because they are negatively charged, add another electron = 24.
Avoid O–O bonds.
Check for formal charges.
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Steps for Building a Dot Structure
Ammonia, NH3
1. Add up the number of valence electrons that can be used.
H = 3 x1 = 3
N = 1x5=5
Total = 8 electrons / 4 pairs
For an ion, add one electron for each negative charge or subtract
one electron for each positive charge on the ion.
Note: Number of valence electrons of Group A elements = Group
number in the periodic table
Steps for Building a Dot Structure
Ammonia, NH3
2. Write down the skeletal arrangement of the atoms and connect
them with a single covalent bond (two dots or one dash).
Decide on the central atom; never H. Why?
If there is a choice, the central atom is atom of lowest affinity for electrons.
(Most of the time, this is the least electronegative atom).
N is central on this one
H
N
H
H
Steps for Building a Dot Structure
Ammonia, NH3
3. Subtract two electrons for each single bond you used in Step 2
from the total number of electron calculated in Step 1.
8 electrons – 6 electrons = 2 etotal # e-
net number of electrons
total e- used to form bonds
This calculation gives you the net number of electrons available for completing
the structure.
Steps for Building a Dot Structure
Ammonia, NH3
4. Distribute pairs of electrons (pairs of dots) around each atom
(except hydrogen) to give each atom a total of eight electrons
around it.
Remaining electrons form LONE PAIRS to complete the octet as needed
(or duet in the case of H).
3 BOND PAIRS and 1 LONE PAIR.
Note: that N has a share in 4 pairs (8 electrons), while H shares 1 pair.
••
H
N H
H
Steps for Building a Dot Structure
Ammonia, NH3
5. If there are not enough electrons to give these atoms eight
electrons, change single bonds between atoms to double or triple
bonds by shifting unbonded pairs of electrons as needed.
Check to see that each atom (except H) has eight electrons around it. A
double bond counts as four electrons for each atom to which it is
bonded.
Excess or remaining electrons are distributed to the central atom (manyelectron atom: more than an octet)
Steps for Building a Dot Structure
Ammonia, NH3
6. Assign formal charges.
Formal charge is the charge an atom would have if all of the electrons in a
covalent bond were shared equally.
predominant
Writing Lewis Structures
The dominant Lewis structure is:
• the one in which atoms have formal charges closest to zero.
• puts a negative formal charge on the most electronegative atom
(F > O > N)
As such, it can be used to decide which structure is best.
NCS-
Problems
Draw the Lewis structure of the following:
1)
2)
3)
4)
5)
PCl3
O2
N2
CO2
HCN
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Kekulé Structures and
Condensed Structures
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What is an Atomic Orbital?
An atomic orbital is a three-dimensional region around the nucleus where an electron is
most likely to be found.
s orbital is a sphere with the nucleus at its center.
The average distance from the nucleus is greater for an electron in a 2s orbital than it is for
an electron in a 1s orbital.
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A Standing Wave
Electrons have both particle-like and wave-like properties. A node is a consequence of the
wave-like properties of an electron.
The vibrating string of a guitar is a standing wave—the string moves up and down, but does
not travel through space.
There is zero probability of finding an electron at the node.
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The Lobes of a p Orbital Have
Opposite Phases
p orbitals have two lobes. They are generally depicted as teardrop shaped.
Computer-generated representations reveal that they are more like doorknobs.
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Orbital node (node): A point or plane of zero electron density in an orbital. Always bordered
by two or more orbital lobes.
The Three p Orbitals
Three p orbitals have the same energy.
Each p orbital is perpendicular to the other two p orbitals.
The energy of a 2p orbital is slightly greater than that of a 2s orbital
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Molecular Orbital Theory
Molecular orbital (MO) theory combines the tendency of
atoms to fill their octets by sharing electrons (the Lewis
model) with their wavelike properties, assigning electrons
to a volume of space called an orbital.
According to MO theory, covalent bonds result when
atomic orbitals combine to form molecular orbitals .
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The Bonding in H2
How do atoms form covalent bonds in order to form molecules?
The covalent bond is formed when the 1s orbital of one hydrogen atom overlaps the 1s
orbital of a second hydrogen atom.
The covalent bond that is formed when the two orbitals overlap is called a sigma (σ) bond.
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The covalently bonded
atoms are more stable
than the individual
atoms
Energy
increases due
to repulsion
Energy released
when the bond is
formed.
Energy required to
break the bond.
Maximum stability
(minimum energy) is
achieved.
Bond length of the new
covalent bond
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Waves Can Reinforce Each Other;
Waves Can Cancel Each Other
Orbitals are conserved . The number of molecular orbitals formed must equal the number
of atomic orbitals combined.
In describing the formation of an H-H bond, we combined two atomic orbitals but
discussed only one molecular orbital. Where is the other molecular orbital?
s bonding MO
s* antibonding MO
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Atomic Orbitals Combine
to Form Molecular Orbitals
(Molecular Orbital Theory)
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Side-to-Side Overlap of In-Phase
p Orbitals Forms a π Bond
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The p Orbital of the More Electronegative
Atom Contributes More to the Bonding MO
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Overlap and Bonding
We think of covalent bonds forming through the
sharing of electrons by adjacent atoms.
In such an approach this can only occur when
orbitals on the two atoms overlap.
Pi () Bonds
Pi bonds are
characterized by
Side-to-side overlap.
Electron density above
and below the
internuclear axis.
The Bonding in Methane
Not partially charged atoms, therefore there is a similar electronegativities between carbon
and hydrogen, which cause them to share their bonding electrons relatively equally.
Methane, therefore, is a non polar molecule.
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In Order to Form Four Bonds,
Carbon Must Promote an Electron
C = 1s22s22p2
Carbon has only two unpaired valence electrons but would not complete its octet if it
will form only two covalent bonds.
The four C¬H bonds in methane are identical because carbon uses hybrid atomic orbitals.
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Four Orbitals are Mixed to Form
Four Hybrid Orbitals
An sp3 orbital has a large lobe and a small lobe.
Hybrid orbitals are
mixed orbitals that
result from
combining atomic
orbitals
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sp3=one part s, three
parts p.
=25% s character,
75% p character.
The Carbon in Methane is sp3
The larger lobe of the sp3 orbital is used to form covalent bonds.
Each of the four C¬H
bonds in methane is
formed
from the overlap of an
sp3 orbital of carbon
with the s orbital of a
hydrogen
Carbon is tetrahedral.
The tetrahedral bond angle is 109.5°.
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The Bonding in Ethane
Each carbon in ethane
(CH3CH3) is bonded to four
other atoms. Thus, both
carbons are tetrahedral.
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The Bonding in Ethane
All the bonds in methane and ethane are sigma (s) bonds. All single bonds in organic
compounds are sigma bonds.
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End-On Overlap of Orbitals
Forms a (σ) Bond
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Bonding in Ethene
Each of the carbon atoms in
ethene forms four bonds, but
each carbon is bonded to only
three atoms.
To bond to three atoms, each
carbon hybridizes three atomic
orbitals: an s orbital and two of
the p orbitals.
three sp2 orbitals
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After hybridization, each carbon atom has three sp2 orbitals and one unhybridized p orbital.
An sp2 Carbon Has Three sp2 Orbitals
and One p Orbital
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The Carbons in Ethene are sp2
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Bonding in Ethyne
Each of the carbon atoms in ethyne forms four bonds, but each carbon is bonded to only
two atoms—a hydrogen and another carbon
Each carbon atom in ethyne, therefore, has two sp orbitals and two unhybridized p orbitals.
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The Two sp Orbitals Point in Opposite Directions
The Two p Orbitals are Perpendicular
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The Carbons in Ethyne are sp
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Methyl Cation and Methyl Radical are sp2
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Methyl Anion is sp3
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Nitrogen Has Three Unpaired Valence Electrons
and Forms Three Bonds in NH3
N = 1s22s22p3
Nitrogen does not have to promote an electron.
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The Bonds in Ammonia (NH3)
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The Ammonium Ion (+NH4)
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Oxygen Has Two Unpaired Valence
Electrons and Forms Two Bonds in H2O
O = 1s22s22p4
Oxygen does not have to promote an electron.
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The Bonds in Water (H2O)
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The Bond in a Hydrogen Halide
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Overlap of an s Orbital with an sp3 Orbital
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The Length and Strength of a
Hydrogen Halide Bond
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Summary of Hybridization
orbitals used in bond formation determine the bond angle
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Single Bond: 1 σ + 1 π
Double bond: 1 σ + 2 π
Triple Bond: 1 σ + 2 π
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Hybridization of C, N, and O
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Bond Strength and Bond Length
The shorter the bond, the stronger it is.
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s Character
The shorter the bond, the stronger it is.
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The More s Character in the Orbital,
the Shorter and Stronger the Bond
The more s character, the shorter and stronger the bond.
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The More s Character in the Orbital,
the Greater the Bond Angle
The more s character, the greater the bond angle.
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Hybridization, Bond Angle,
Bond Length, Bond Strength
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Summary
• The shorter the bond, the stronger it is.
• The greater the electron density in the region of orbital
overlap, the stronger the bond.
• The more s character, the shorter and stronger the bond.
• The more s character, the larger the bond angle.
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A π Bond is Weaker Than a σ Bond
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Dipole Moments of Molecules
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Dipole Moments of Molecules
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