Ch01 lecture 01_05

8/27/19
1
1
Chapter 1
Structure Determines Properties
©Science and Society/SuperStock
©2020 McGraw-Hill Education. All rights reserved. Authorized only for instructor use in the classroom. No reproduction or further distribution
permitted without the prior written consent of McGraw-Hill Education.
2©2020 McGraw-Hill Education.
Section 1.1
ATOMS, ELECTRONS, AND
ORBITALS

8/27/19
2
Particles and Symbols of the Atom
The number of protons in the nucleus is called the atomic number (Z).
Electrons as Waves
• Electrons in atoms and molecules behave as waves rather
than particles
ˆH EY = Y The Schrödinger equation
• The wavelike behavior of electrons is captured by the
wavefunction (ψ) or orbital
• Every electron has an associated orbital; the electron occupies the
orbital
• The shape of an orbital reflects the probability
density of the electron over space
• The energy of an orbital reflects the stability
of an electron within it
• The probability density p(r) at a point r is
related to
2
Y

8/27/19
3
Quantum Numbers
• Atomic orbitals are associated with quantum numbers that
characterize the energy and shape of the orbital
• Principal quantum number (n): related to the energy of the
orbital
• Orbital quantum number (l): related to the shape of the
orbital
• Magnetic quantum number (ml): related to the orientation
and directionality of the orbital
• Spin quantum number (s): related to the magnetic
properties of the electron
The Pauli exclusion principle
states that no two electrons in an
atom or molecule can share the
same set of quantum numbers.
The s Atomic Orbitals
• The s atomic orbitals begin at n = 1 and are spherically shaped
• There is a single orbital within each ns subshell
• Within a particular shell, s orbitals are lower in energy than other orbitals
• As n increases, additional sign changes (nodes) appear in the shape of
the orbital

8/27/19
4
The p Atomic Orbitals
The p atomic orbitals begin at n = 2 and are dumbbell shaped
There are 3 orbitals within each np subshell, which correspond to the
three Cartesian directions
The np orbitals all have the same energy; according to Hund’s rule
electrons are left unpaired when filling until they must be paired
All p orbitals contain a node at the nucleus
Atomic Electron Configurations
Rule 1: Orbitals fill in order of
increasing energy from lowest to
highest (1s < 2s < 2p < 3s < 3p)
Rule 2: No 2 electrons can have
the same set of 4 quantum
numbers.
Rule 3: For degenerate orbitals,
each orbital is singly occupied
before it is doubly occupied.

8/27/19
5
Atomic Electron Configurations
TABLE 1.1 Electron Configurations of the First Twelve Elements of the Periodic Table
Element Atomic
number Z
Number of
electrons in
indicated
orbital
1s
Number of
electrons in
indicated
orbital
2s
Number of
electrons in
indicated
orbital
2px
Number of
electrons in
indicated
orbital
2py
Number of
electrons in
indicated
orbital
2pz
Number of
electrons in
indicated
orbital
3s
Hydrogen 1 1
Helium 2 2
Lithium 3 2 1
Beryllium 4 2 2
Boron 5 2 2 1
Carbon 6 2 2 1 1
Nitrogen 7 2 2 1 1 1
Oxygen 8 2 2 2 1 1
Fluorine 9 2 2 2 2 1
Neon 10 2 2 2 2 2
Sodium 11 2 2 2 2 2 1
Magnesium 12 2 2 2 2 2 2
Section 1.2
IONIC BONDS

8/27/19
6
Coulomb’s Law and Ionic Bonding
Ions are atoms or molecules with electric charge.
Oppositely charged ions attract one another
according to Coulomb’s law 1 2q q
E k
r
=
The resulting attraction is known as an ionic bond
Ions pack into tightly clustered lattices in ionic
compounds
• High-melting solids
• Soluble in polar solvents
Ionic Bonding and Electron Transfer
• One can imagine ionic bonds as arising from the transfer of
an electron from a metal to a nonmetal
• For example, gaseous sodium can transfer an electron to
gaseous chlorine…
• This process is endothermic (ΔH degree = + 147 kJ/mol),
but the formation of solid NaCl is strongly exothermic
Gaseous Na and
Cl spontaneously
form solid NaCl.

8/27/19
7
Ionic Bonding in Organic Chemistry
• Carbon atoms rarely form ions; hence, ionic bonds are rare
in organic chemistry
• Covalent bonds involving the sharing of electrons between
nonmetal atoms are much more common
• Ionic bonds do appear in salts of C, N, O, and H which are
nucleophilic (electron-donating) at the nonmetal atom
H3C O Na H3C
C
Li
Section 1.3
COVALENT BONDS, LEWIS
FORMULAS, AND THE OCTET
RULE

8/27/19
8
The Lewis Model of Covalent Bonding
• Covalent bonds involve the sharing of electrons between
two atoms
• Atoms share electrons to achieve a more stable electron
configuration
• Maximum stability is reached when an atom achieves a full
valence shell, isoelectronic with the nearest noble gas
• Electrons may not be shared equally—covalent
bonds may be polarized!
Covalent Bonding in H2
• A neutral hydrogen atom needs one more valence electron
to achieve a full valence shell
• Two hydrogens sharing their valence electrons achieve a full
n = 1 shell
• In H2, the electron configuration of each hydrogen atom is
analogous to that of helium, the first noble gas
Two dots or a line between two atoms denotes the sharing of
two electrons between the atoms in a covalent bond.

8/27/19
9
Covalent Bonding in F2
• A neutral fluorine atom needs one more valence electron to
achieve a full valence shell
• Two hydrogens sharing their valence electrons achieve a
full n = 2 shell containing 8 electrons
• In F2, the electron configuration of each fluorine atom is
analogous to that of neon, the second noble gas
The six electrons on the periphery of each fluorine are
nonbonding electron pairs or lone pairs.
Lewis Symbols and Structures
• The Lewis structure of a molecule represents the arrangement
of valence electrons of the atoms in the molecule (named after
G.N Lewis).
• Electrons are represented as dots or lines. A line denotes a pair
of electrons (2) shared in a covalent bond; atomic symbols
denote atoms + core e−’s.
• Not all electrons will be involved in covalent bonding; unshared
pairs are drawn as dots on the edges of their associated atoms.
• Key premise: electrons are localized on a single atom or
between two atoms.

8/27/19
10
The Octet Rule
• For second-row elements, a full valence shell contains
electrons in the 1s, 2px, 2py, and 2pz orbitals
2 2 2 2+ + + = 8
• A full n = 2 valence shell is called an octet because it
contains eight electrons
• The octet rule: in stable molecules, second-row atoms share
electrons until they achieve an octet
• When checking that the octet rule is satisfied, double-count
shared electrons (once for each atom involved in sharing)
Each fluorine has an
octet of electrons.
Lewis Structures
• In most stable compounds, the atoms achieve noble
gas electron configurations.
• Rules
1. Sum the valence electrons from all the atoms.
• Add 1 electron for each negative charge and subtract 1 electron for
each positive charge.
2. Use electron pairs to form a bond between each pair of
bound atoms.
3. Arrange the remaining electrons to satisfy the duet rule
for hydrogen and the octet rule for the second-row
elements.

8/27/19
11
Lewis Structures
• Generalizations:
¨ H has one bond
¨ C has four bonds
¨ N has three bonds and one unshared pair of electrons
¨ O has two bonds and two unshared pairs of electrons
¨ F, Cl, Br and I have one bond and three unshared pairs of
electrons
• Exercise: Draw the Lewis structures of CO2, HCN, C2H4,
C2H2, H2O2, N2H4.
Examples of Lewis Structures

8/27/19
12
Multiple Bonds
• More than two electrons can be shared between atoms.
This is often necessary to satisfy the octet rule!
• When four or six electrons are shared, a multiple bond
results
• Double bond: four electrons are shared (2 × 2)
• Triple bond: six electrons are shared (3 × 2)
Double Bonds
• To draw the Lewis structure of ethylene (C2H4), we must
share four electrons between the carbon atoms to
achieve octets on both

8/27/19
13
Triple Bonds
• To draw the Lewis structure of ethyne (C2H2), we must
share six electrons between the carbon atoms to achieve
octets on both
Carbon Dioxide Lewis Structure
• To draw the Lewis structure of carbon dioxide (CO2), we
must share four electrons between the carbon and
oxygen atoms to achieve octets on both

8/27/19
14
Section 1.4
POLAR COVALENT BONDS,
ELECTRONEGATIVITY, AND
BOND DIPOLES
Electronegativity
• Electronegativity is defined as the ability of an atom to
attract electrons to itself
• Electronegative atoms attract electrons strongly, hold their
electrons tightly, and tend to take on electrons
• Electropositive atoms attract electrons weakly and may give
up electrons
Electronegativity helps us predict the
relative reactivity of analogous
compounds. Recalling trends in
electronegativity is extremely important!

8/27/19
15
Periodic Trends in Electronegativity
TABLE 1.3 Selected Values from the Pauling Electronegativity Scale
Period
Group
number
1A
Group
number
2A
Group
number
3A
Group
number
4A
Group
number
5A
Group
number
6A
Group
number
7A
1
H
2.1
2
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
0
3.5
F
4.0
3
Na
0.9
Mg
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
4
K
0.8
Ca
1.0
Br
2.8
5
l
2.5
Electronegativity increases from left to right across a period and
from bottom to top within a group.
Polar Covalent Bonds
• Two atoms of different electronegativities share electrons
unequally in a covalent bond. The result is a polar covalent
bond
• The greater the difference in electronegativity, the more
polarized the bond
Polar covalent bonds tend to be sites of
reactivity in organic molecules.

8/27/19
16
Electrostatic Potential Maps1
• The model of molecular charge as a dipole is a simplification;
In reality molecules contain a spatial distribution of charge
• An electrostatic potential (ESP) map shows the distribution
of charge over a molecule
Electrostatic Potential Maps2
• The model of molecular charge as a dipole is a simplification;
In reality molecules contain a spatial distribution of charge
• An electrostatic potential (ESP) map shows the distribution
of charge over a molecule

8/27/19
17
A Selection of Dipole Moments
More polarized bonds are associated with greater differences
in electronegativity and greater dipole moments.
Section 1.5
FORMAL CHARGE

8/27/19
18
Formal Charge
• Within a Lewis structure, an atom may not have a formal
number of electrons equal to the valence electron count of
the neutral atom
• In this case, the atom has nonzero formal charge
• To calculate formal charge, subtract the valence electron
count of the neutral atom (VEC) from the formal valence
electron count of the covalently bound atom (FEC)
FC FEC VEC= -
Formal charges with magnitude greater than
1 are not encountered in organic molecules.
Calculating Formal Charge1
• To calculate the formal valence electron count of an atom,
count all unshared electrons and half of all bonding electrons
• Refer to the periodic table (group number) for the number of
valence electrons in the neutral atom
A neutral oxygen atom has 6
valence electrons; the formal
charge of this O is thus 6 − 7 = −1.

8/27/19
19
Calculating Formal Charge2
• Stay on the lookout for structural patterns associated with
formal charge, such as nitrogen with four bonds and oxygen
with one bond
A neutral nitrogen atom has 5
charge of N is thus 5 − 4 = +1.
A neutral oxygen atom has 6
charge of this O is thus 6 − 7 = −1.
Formal Charge in NH4
+
• A neutral nitrogen atom contains 5 electrons (Group 5A)
• The nitrogen atom in NH4 has a formal valence electron
count of 0.5 × 8 = 4
• The formal charge of nitrogen is thus 5 − 4 = +1

Ch01 lecture 01_05

Recommended

Recommended

More Related Content

What's hot

What's hot (20)

Similar to Ch01 lecture 01_05

Similar to Ch01 lecture 01_05 (20)

Recently uploaded

Recently uploaded (20)

Ch01 lecture 01_05