Transcript Types of Chemical Bonds

Types of chemical bonds
Bond: Force that holds groups of two or more atoms
together and makes the atoms function as a unit.
Example:
H-O-H
Bond Energy: Energy required to break a bond.
Ionic Bond: Attractions between oppositely charged ions.
Example:
Na+ Cl-
Types of chemical bonds
Ionic Compound: A compound resulting from a positive ion
(usually a metal) combining with a negative ion (usually
a non-metal).
Example:
M+ + X-  MX
Covalent Bond: Electrons are shared by nuclei.
Example:
H-H
Polar Covalent Bond: Unequal sharing of electrons by
nuclei.
Example:
H-F
Hydrogen fluoride is an example of a molecule that has
bond polarity.
Lewis structures
Lewis Structure: Representation of a molecule that shows
how the valence electrons are arranged among the atoms
in the molecule.
Bonding involves the valence electrons of atoms.
Example:
Na●
H-H
Lewis structures of
elements

Dots around elemental symbol
– Symbolize valence electrons

Thus, one must know valence electron
configuration
Lewis Structures of
molecules
Single Bond: Two atoms sharing one electron pair.
Example: H2
Double Bond: Two atoms sharing two pairs of electrons.
Example: O2
Triple Bond: Two atoms sharing three pairs of electrons.
Example: N2
Resonance Structures: More than one Lewis Structure can
be drawn for a molecule.
Example: O3
Rules for Lewis structures of
molecules
1.
2.
Write out valence electrons for each atom
Connect lone electrons because lone
electrons are destabilizing
1. Become two shared electrons
1. Called a “bond”
3.
Check to see if octet rule is satisfied
1. Recall electron configuration resembling noble
gas
1.
In other words, there must be 8 electrons (bonded
or non-bonded) around atom
1. Non-bonded electron-pair
1.
Called “lone pair”
Let’s do some examples
on the board
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H2
– Duet rule

F2
– Octet rule
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O2
N2
Lewis structures
Example
Write the Lewis Structure for the following
molecules:
1)
H2O
2)
CCl4
1) Where does the carbon go & why?
3)
4)
5)
PH3
H2Se
C2H6
Lewis structures
continued
6)
7)
8)
9)
CO2
C2H4
C2H2
SiO2
Polyatomic ions

If positive charge on ion


If negative charge on ion


Take away electron from central species
Add electron to central species
Example:

H3O+
Your turn
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NH4+
ClOOH-
Resonance structures

When structures can be written in
more than one way
– O3

Actual molecule is “in-between”
– Resonance hybrid

Another example
– HCO3
What would its resonance hybrid look like?
Practice
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NO2NO3-
Formal Charge
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1.
2.
3.
Charge calculated on atom based on Lewis
structure
– Yields best Lewis structure of competitors
FC = VE - [LE + ½(BE)]
Rules:
Sum of all FC’s must equal to charge on
species, if any
Smaller or zero FC’s on atoms better than
large FC’s
Negative FC should be on most
electronegative species
Examples

HBr
– FC on H = 1-[0 + ½ (2)] = 0
– FC on Br = 7 – [6+ ½ (2)] = 0
– Net sum of FC’s = charge on ion = 0

OH– FC on O = 6 – [6 + ½(2)] = -1
– FC on H = 1 – [0 + ½(2)] = 0
– Net sum of FC’s = charge on ion = -1
Practice


H2O2
H3O+
Aberrant compounds

Odd-electron species
– NO
– NO2
Aberrant compounds

Incomplete octet
– BH3
Aberrant compounds

Expanded octet
– Some central atoms can exceed an octet

Third period and higher elements can
do this
– E.g., Al, Si, P, S, Cl, As, Br, Xe, etc.
– d-orbitals can accommodate extra
electrons
Examples
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
AsI5
XeF2
Practice
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
SCl6
XeF4
Aberrant compounds

Write this out:
– SO42-

Can we reduce the formal charges?
– If so, how?

We can also find the average FC
– Let’s take a look
Aberrant compounds

Its OK to expand the octet for those
atoms that can take it in order to
lower FC’s
Practice
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
SO32PO33SO2
SO3
H2SeO4
Electroneutrality principle

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Electrons distributed so that charges on
atoms are closest to zero
If “-” charge present, should be on most
electronegative atom
(so, “+” charge should be on least
electronegative atom)
Good for deciding which resonance
structure is best
Example: OCN-
Electronegativity
Electronegativity: The relative ability of an atom in a
molecule to attract shared electrons to itself.
Example: Fluorine has the highest electronegativity.
 Similar electronegativities between elements give nonpolar covalent bonds (0.0-0.4)
 Different electronegativities between elements give
polar covalent bonds (0.5-1.9)
 If the difference between the electronegativities of two
elements is about 2.0 or greater, the bond is ionic
Electronegativity
Example
For each of the following pairs of bonds,
choose the bond that will be more polar.
 Al-P vs. Al-N
 C-O vs. C-S
Dipole moment


Dipole Moment
 A molecule that has a center of positive
charge and a center of negative charge
 Will line up on electric field
In Debye units
 1 D = 3.34 x 10-30 C  m
Examples
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F2
CO2
H2O
NH3
BF3
CCl4
Molecular polarity
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
Net-dipole moment leads to molecular
polarity
Thus the following two that have netdipole moments are polar:
– H2O
– NH3
Molecular structure
Molecular Structure: or geometric structure refers to the
three-dimensional arrangement of the atoms in a
molecule.
Bond Angle: The angle formed between two bonds in a
molecule.
Molecular structure:
VSEPR
The VSEPR Model: The valence shell electron pair
repulsion model is useful for predicting the molecular
structures of molecules formed from nonmetals.
The structure around a given atom is determined by
minimizing repulsions between electron pairs.
The bonding and nonbonding electron pairs (lone pairs)
around a given atom are positioned as far apart as
possible.
Molecular Structure:
VSEPR
Steps for Predicting Molecular Structure Using the VSEPR
Model
1. Draw the Lewis structure for the molecule.
2. Count the electron pairs and arrange them in the way
that minimizes repulsion (that is, put the lone pairs as
far apart as possible).
3. Determine the positions of the atoms from the way the
electron pairs are shared.
4. Determine the name of the molecular structure from the
positions of the atoms.
Example




Br2
CO2
CF4
PF3
Your turn
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NH4+
XeF4
AsI5
SF3 +
I3 -