Chemical Bonding
1
Objectives
Compare and contrast the types of chemical
bonds , energy involved in their formation,
lengths, angles and their structure.
Main types:
covalent – molecular
ionic
metallic
2
Ionization Energy
• Ionization Energy – energy needed to remove
an e- from the outermost shell of a neutral
atom. (A measure of the ability of an atom or
ion to hold onto electrons.)
– Low i.e. means easier removal of e- and a
resultant positively charge cation will be formed
– Trend: greater i.e. up and to the right on the
periodic chart
3
Ionization Energy
• A measure of the ability of an atom or ion to
hold onto electrons.
• Trend: i.e. increases up and to the right.
– Electrons are held more tightly by positive ions.
– Electrons are held less tightly by negative ions.
4
Electron Affinity
• The energy released or absorbed by a neutral
atom from the acquisition of an electron to its
outer shell.
– High e.a. means it’s easier to accept an e- and a
resultant negatively charged anion will be formed.
– Trend: highest e.a. at top of group because of
increased nuclear attraction for acquired e- and
highest at right of each period because atoms
reach stability by achieving stability because of the
acquisition
5
Types of Chemical Bonds
1. Covalent – electrons are shared
a) Non-polar – e’s shared equally
2.
3.
4.
5.
6.
Polar covalent – unequal sharing
Ionic – electrons are transferred
Coordinate covalent
Radical
Metallic
6
Electronegativity
A measure of the ability of an atom in a chemical
compound to attract electrons. EN can be used to
determine bond type.
7
1. Non-Polar Covalent Bonds
• These occur primarily between two nonmetallic elements, especially the diatomic
gases (H, N, O, and the Halogen family)
• There is an equal sharing of the valence
electrons so that both atoms achieve octet
and chemical stability
• e. n. falls between 0.0 and 0.4
Source:
http://www.attanolearn.com/excel/4317_co
valent-bond.jsf
8
2. Polar Covalent Bonds
• These occur primarily between two active
non-metals or between a moderately active
metal and a non-metal. There is an unequal
sharing of the valence electrons.
• e.n. difference is 0.5 to 1.6 range.
Electrons are held
closer to the Oxygen
because O has greater
e.n.
Source: academic.brooklyn.cuny.edu
9
Another way to show polar covalent
d means slightly – or +
Source: bioactive.mrkirkscience.com
10
Main Bond Types
source: homework-help-secrets.com
11
3. Ionic Bonds
• These occur primarily between active metals
and active non-metals. The lesser
electronegative element actually transfers one
or more valence electrons to another atom –
the more e.n. element.
• Range of e.n. difference is 1.7 to 4.0.
• See prior slide for sample.
12
Ionic Bond Sample
•
Source: kentsimmons.uwinnipeg.ca
13
Lattice Energy
The energy released when two elements combine during
the formation of a compound due to electrostatic interactions
forming the molecule’s physical structure.
Ions of unlike charge are attracted to one another.
Equilibrium is reached and a lattice is formed (as in the
picture at upper right).
Energy released in formation of the compound is equal to
the energy needed to break apart the compound into its
component elements.
Source:
http://www.science.uwaterloo.ca/~cchi
eh/cact/c120/chembond.html
14
4. Coordinate Covalent Bonds
forming Polyatomic Ions
• These occur primarily between two nonmetals one of which is usually oxygen. The
lesser e.n. element provides both shared
electrons.
• Always results in the formation of a
polyatomic ion – usually an anion or radical.
nitrate ion, NO3 1-
phosphate ion, PO4 315
5. Radical Bonds
• Occur mainly between a metallic cation and a
radical anion. Cation transfers e- to the central
atom of the radical which shares these
electrons with its combining atoms by
coordinate covalent bonds.
• Contains both ionic and covalent bonds.
• Always forms a stable multi-atomic compound
of at least 3 different elements.
16
Radical Bond Samples
Sodium Phosphate
Na3(PO4)
Sodium Nitrate
NaNO3
17
6. Metallic Bonds
Occur only between metals during the
formation of alloys. The metallic kernel,
composed of metallic nuclei and their inner
shell electrons, is surrounded by a “sea” of
valence electrons that flow across and about
the kernel.
Metal cation
- electrons that are delocalized
or free to move about
Source: www4.nau.edu
18
• Metals are shiny because they absorb light,
exciting electrons to higher energy levels. The
e’s immediately fall to lower levels, emitting
light energy and making the metals appear
shiny.
• Ductile
• Malleable
19
Electronegativity Chart
20
Relationship Between Electronegativity
Difference and Ionic Character
Electronegativity
Percentage of
Difference
Ionic Character
0.2 nonpolar
1
0.4 covalent
4
bond
0.5
0.6
9
0.8 polar
15
1.0 covalent
22
1.2 bond
30
1.4
39
1.6
47
1.8
55
2.0
63
2.2
70
2.4 ionic
76
2.6 bond
82
2.8
86
3.0
89
3.2
92
(See page 4 of note handout
or page 161 in book or prior slide)
21
Sample e.n. problems
See pages 4, 5, & 6 of note handout
or frames 20 and 21
Given: As2S3
e.n. difference:
2.44 – 2.20 = 0.24
Bond Type:
Non-polar covalent
Given: CaF2
e.n. diff.:
4.10 – 1.04 = 3.06
Bond Type:
Ionic
22
Assignment
• Ch 6 Review pages 41 & 42
• #30 177/1-6
6.1
• #33 209/1-7
6.1
23
Covalent Bond Characteristics
• Atoms bond to become more stable by getting
a full outer energy level (Octet for all except H
and He).
24
From table at left notice
that shorter bond lengths
require a greater energy to
break the bond and form
neutral isolated atoms.
Notice how in the figure at
left how a bond is formed
between 2 hydrogen atoms
to form a stable 1s2
configuration of H2.
Source: Modern Chemistry, 2006 ed.
25
There are exceptions to
the octet rule: some
compounds formed
with F, O, and Cl. This
is called expanded
valence.
Source: Modern Chemistry, 2006 ed.
26
Electron-Dot Notation
Dots are placed around the
symbol of an element to
represent its number of
valence electrons.
Source: Modern Chemistry, 2006 ed.
27
Lewis Structures
• Formulas in which atomic symbols represent
nuclei and inner-shell electrons, dot-pairs or
dashes represent electron pairs in covalent
bonds. Adjacent dots represent unshared
electrons. (Structural formulas show dashes
only: F-F or H-Cl)
28
Drawing a Lewis structure
1. Determine type and number of atoms in the
molecule
2. Write electron-dot notation for each atom
type
3. Determine total number of valence electrons
4. Arrange atoms. C is central atom if present;
or least e-n atom is central (not H).
5. Arrange electrons to get 8 around each atom
(except H)
29
NH3
CO2
30
NH3
CO2
Source: imagesfrom.us
Source: mcat-review.org
31
Assignment
•
•
•
•
#30
#34
#35
#31
189/1-5
209/8-15
209/16-24
194/1-5 and 196/1-3
6.2
6.1&2
6.2
6.3&4
32
Resonance Structures
Molecules or ions that cannot be correctly represented
by a single Lewis structure. Below are three possibilities
for NO3-1 - the nitrate ion.
www.nku.edu/~russellk/tutorial/reson/resonance.html
33
Ionic Bonding & Ionic Compounds
• Formed by electrons being transferred. The
number of positive and negative ions are
equal (no charge). Simplest form is a formula
unit.
NaCl
MgO
CaCl2
34
Info on Ionic Bonds
• They form a crystal lattice which minimizes
their potential energy.
• Stronger bonds than in covalent compounds
producing higher boiling or melting points
than cov.
• Hard but brittle.
• Do not conduct in solid state, but do conduct
in molten state or when dissolved in water.
35
Hybridization and Molecular Geometry
• VSEPR (Valence-shell, electron-pair repulsion) theory is
used to predict shapes of molecules based on the fact
that unshared electron pairs strongly repel each other
• Hybridization theory is used to predict the shapes of
molecules based on the fact that orbitals within an
atom can mix to form orbitals of equal energy.
• Bottom line: bonding in atoms depends on the central
atom’s ability to maximize the spread of its valence
electrons into adjacent orbitals within the same energy
level.
36
You have this sheet (p. 7) in
your Ch 6 handout…
It gives you compound
samples such as BeF2 and BF3
and shows you electron pairs
that are shared and unshared,
hybrid type, angles and
geometry.
37
Page 7 of 10 in handout
38
Page 9 of your Ch6 handout:
Molecular Type – AXE format
A = central atom
X = shared pairs of electrons
(shown as B in your book)
E = unshared pairs of electrons
Shared pairs = number of
bonding atoms united with the
central atom of the molecule
3 D Diagram meanings
shared pairs
above, below or to the
side of the central
atom
behind of central atom
in front of central atom
..
lone pairs of electrons
39
Dipole-Dipole Forces
So by equal but opposite
• A dipole is created
ur
ce
forces that are separated
by a short distance
:
H
H
H
Hf
f.g
pc
.e
du
Source: facstaff.gpc.edu
Hydrogen bonding – between H of one
molecule of water and O of another water molecule. Shown by
dashed line.
40
• A polar molecule can induce a dipole in a
nonpolar molecule by temporarily attracting
its electrons. Making O2 dissolve in H20, for
instance. See page 206 in book.
• London dispersion forces cause intermolecular
attractions as a result of constant motion of
electrons and creation of instantaneous
dipoles.
41
Assignments
•
•
•
•
#32
#36
#37
#38
207/1-5
210/25-32
210/33-42
211/43-49
6.5
6.3&4
6.5
Practice
• Chem review from Prentice Hall.
http://wps.prenhall.com/esm_mcmurry_funda
mentals_4/38/9913/2537971.cw/index.html
42