Chapter 8 Chemical Bonding

Transcription

Chapter 8
Chemical Bonding
•
•
•
•
Types of Bonds
Ionic Bonding
Covalent Bonding
Shapes of Molecules
8-1
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Table 8.1 Two Carbon Compounds
Compound
Calcium
Carbonate
Carbon Dioxide
Formula
CaCO3
CO2
Physical State
White solid
Colorless gas
Molar Mass (g/mol)
100.1
44.01
Density (g/mL)
2.71
0.00198
Melting Point (°C)
1339 (at high P)
-56.6 (at 5.11 atm)
Boiling Point (°C)
Decomposes
Sublimes at 78.6
Electrical Conductivity
as a Liquid
High
Very low
Dissolves in
Acids
H2O, CCl4
8-2
Table 8.2 General Properties of Ionic and
Covalent Compounds
Ionic
Covalent
Crystalline solids
Gases, liquids, or solids
Hard, brittle solids
Weak, brittle solids or soft
and waxy solids
Very high melting point
Low melting point
Very high boiling point
Low boiling point
Good electrical conductor
when molten or in solution
Often soluble in water but not
in carbon tetrachloride
Poor conductor of electricity
and heat
Often soluble in carbon
tetrachloride but not in water
8-3
1
Chemical Bonds
• Chemical bond
– A force that holds
atoms together in a
molecule or
compound
Figure 8.2
• Two types of
chemical bonds
– Ionic Bonds
– Covalent Bonds
8-4
Ionic Bond
•
•
•
•
A bond created by
electrostatic attraction
between oppositely
charged ions
Occurs between a metal
and a nonmetal
Electrons transferred
between the cation
(positively charged ion)
and the anion (negatively
charged ion)
Extremely strong bonds
8-5
Covalent Bonds
Figure 8.2 or another
•
molecule picture
•
•
•
•
A bond created by the
sharing of electrons
between atoms
Occurs between two
nonmetals (resulting in
a neutral overall
charge)
Electrons not
transferred in this case
Electrons shared in
pairs typically
Weaker bonds than
ionic bonds
8-6
2
Practice – Identifying Types of
Bonding
• Identify the type of bonding in each
of the following substances:
1. NaF
2. ClO2
3. FeSO4
4. SO2
5. Ca(ClO2)2
8-7
Practice Solutions – Identifying
Types of Bonding
• Identify the type of bonding in each of the
following substances:
1. NaF – Ionic bonding (metal + nonmetal)
2. ClO2 – Covalent bonding (2 nonmetals)
3. FeSO4 – Ionic bonding between the metal
and nonmetal; Covalent bonding between
the nonmetals in the polyatomic ion
4. SO2 – Covalent bonding
5. Ca(ClO2)2 - Ionic bonding between the metal
and nonmetal; Covalent bonding between
the nonmetals in the polyatomic ion
8-8
Polar vs. Nonpolar
• Two general types of covalent bonds:
– Polar covalent
• Unequal sharing (or a partial transfer) of
electrons
• Occurs when different elements are
covalently bonded to one another
–Why different elements?
»Because different elements have
different electronegativities
– Nonpolar covalent
• Equal sharing (no transfer) of electrons
• Occurs only when all of the atoms in a
molecule belong to the same element
8-9
3
Polar vs. Nonpolar (Figure 8.4)
Bonds
Ionic
Covalent
Polar
Complete Transfer
of Electrons
No Sharing
of Electrons
Nonpolar
Increasing electron
transfer
Increasing equality of
sharing
No Transfer
of Electrons
Equal Sharing
of Electrons
8 - 10
Polar vs. Nonpolar
• Polar covalent bonds are:
– Typically shorter bonds
– Stronger bonds due to their increased ionic
character
• Nonpolar covalent bonds are:
– Typically longer bonds
– Weaker bonds
• Polarity
– Occurs in polar covalent molecules
– Polarity is the degree of transfer of electrons
in a covalently bonded molecule composed
of different element’s atoms.
8 - 11
Electronegativity
•
•
•
Ability of an atom to attract bonding electrons
Proposed by Linus Pauling in the early 1930’s
A difference in electronegativity between the
atoms in a covalent bond results in:
– A polar covalent bond
– Increased ionic character
• The greater the difference in electronegativity,
the greater the ionic character and the more polar
the bond that joins the atoms.
– Decreased bond length and increased bond
strength
•
No difference in electronegativity between
atoms in a covalent bond results in a nonpolar
covalent bond.
8 - 12
4
Electronegativity
Figure 8.5
8 - 13
Trends in Electronegativity
Figure 8.6
8 - 14
Practice – Polar Bonds
• Which of the following molecules
have polar bonds? If a bond is
polar, which atom has a partial
negative charge?
1. SO2
2. N2
3. PH3
4. CCl4
5. O3
8 - 15
5
Practice Solutions – Polar Bonds
•
Which of the following molecules have polar
bonds? If a bond is polar, which atom has
partial negative charge?
1. SO2 – Polar covalent bonds O is more
electronegative and has a partial negative
charge
2. N2 – Nonpolar covalent bonds
3. PH3 – Polar covalent bonds P is more
charge
4. CCl4 – Polar covalent bonds Cl is more
charge
5. O3 – Nonpolar covalent bonds
8 - 16
Ionic Bonding
Formation of ions and ionic bonds relates to an
element’s electron configuration.
• Each element immediately following a noble
gas is a metal.
– Metals lose electrons, forming a positive
charge, to become cations.
• Each element immediately preceding a noble
gas is a nonmetal.
– Nonmetals gain electrons, forming a
negative charge, to become anions.
• Therefore, elements (main-group) either lose or
gain electrons to become isoelectronic with a
noble gas (i.e. have the same electron
configuration).
•
8 - 17
Ionic Bonding
8 - 18
6
Lewis Dot Symbols
• Lewis Dot symbol
– Electron dot symbol
– Dots placed around an element’s symbol
represent valence electrons
– Pair electrons as needed
– Octet rule
• Tendency of an atom to achieve an electron
configuration having 8 valence electrons
– Same as the electron configuration of a
noble gas
– The 8 electrons exist in 4 pairs
– Ions achieve 8 electrons by losing or
gaining electrons
8 - 19
Practice – Lewis Symbols for Ions
• Write the Lewis symbols for the
beryllium and nitrogen ions.
Then write a formula for the
compound that would form
between them, using their Lewis
symbols.
8 - 20
Practice Solutions – Lewis
Symbols for Ions
•
Write the Lewis symbols for the beryllium and nitrogen
ions. Then write a formula for the compound that
would form between them, using their Lewis symbols.
Lewis Symbols for beryllium and nitrogen.
N
Be
N
Be
Be
N
Be
N
N
Be
2
N
2
3
3
3
N
Lewis Structures
for beryllium and
nitrogen ions.
Compound
formula
Be3N2
8 - 21
7
Structures of Ionic Crystals
• Crystal lattice
– The pattern obtained when an ion, represented as a
charged sphere, exerts a force equally in all
directions.
• Thus, ions of equal and opposite charge surround it.
– Cations and anions must come into contact for a
crystal lattice to form.
Figure 8.10
8 - 22
• Ionic crystal
– Ions are arranged in a regular geometric pattern that
maximizes the attractive forces and minimizes the
repulsive forces.
– Hard and brittle
– Can shatter if struck forcefully
• The charges and sizes of ions largely determine the
characteristic patterns of ionic crystals
8 - 23
8 - 24
8
Octet Rule
•
Octet rule
– Tendency of an atom to achieve an electron
configuration having 8 valence electrons
• Same as the electron configuration of
a noble gas
• Covalently bonded atoms achieve 8
valence electrons by sharing electrons
• The 8 electrons exist in 4 pairs
– H reacts to obtain a total of 2 electrons
like He.
8 - 25
Covalent Bonding
•
Single covalent bond
–
–
–
A covalent bond that consists
of a pair of electrons shared
by two atoms
Each atom contributes one
electron to the bond
• The orbitals overlap to
allow the electron pair to
be located around both
atoms
Lewis formula
• The atoms are shown
separately and the valence
electrons are represented
by dots
Figure 8.14
8 - 26
Covalent Bonding
•
Multiple covalent bonds
– Covalent bonds that consist
of more than one pair of
electrons shared by two
atoms
– Double bond
•
•
Figure 8.17
Sharing of two pairs of
electrons (4 electrons total)
In Lewis Dot structures, a
double bond is represented by
4 dots or 2 parallel lines.
– Triple bond
•
•
Sharing of three pairs of
electrons (6 electrons total)
In Lewis Dot structures, a
triple bond is represented by 6
dots or 3 parallel lines.
8 - 27
9
Practice – Lewis Formulas
• Determine the formula
of a simple compound
that follows the octet
rule and is formed
from nitrogen and
fluorine atoms. Use
electron dot structures
to describe the
bonding in this
compound.
Figure 8.18
8 - 28
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Practice Solutions – Lewis Formulas
•
Determine the formula of a simple compound that
follows the octet rule and is formed from nitrogen and
fluorine atoms. Use electron dot structures to describe
the bonding in this compound.
Lewis Symbols for nitrogen and fluorine.
F
N
F
N
F
F
FNF
F
F N F
F
8 - 29
Steps for Writing Lewis
Dot Structures
1. Write an atomic skeleton.
•
•
•
•
•
The arrangement of atoms is usually symmetrical.
In a molecule of two different elements, the one with
the greater number of atoms usually surrounds the
one with the lesser number of atoms.
The central atom, the one surrounded by the other
atoms, tends to be the one that is less
electronegative and is present in the least quantity.
This atom usually forms the greater number of bonds
and is found further toward the bottom left side of
the periodic table.
Hydrogen atoms are generally on the outside of the
molecule.
The chemical formula may give clues about the
arrangement of atoms.
8 - 30
10
Dot Structures (Cont’d)
2. Sum the valence electrons from each atom to
get the total number of valence electrons.
3. Place two electrons, a single bond, between
each pair of bonded atoms.
4. If you have not placed all the valence electrons
in the formula, add any remaining electrons as
unshared electron pairs, consistent with the
octet rule.
•
Add pairs of electrons first to complete the octet of
atoms surrounding the central atom. Then add any
remaining electrons in pairs to the central atom.
8 - 31
Dot Structures (Cont’d)
5. If necessary to satisfy the octet rule,
shift unshared electrons from nonbonded position on atoms with
completed octets to positions
between atoms to make double or
triple bonds.
8 - 32
Writing Lewis Dot Structures
• Write a Lewis formula for the
formaldehyde, CH2O, molecule.
CH2O
• Write a Lewis formula for cyanic
acid.
HCN
8 - 33
11
1. Write an atomic skeleton for CH2O:
O
C
H
H
2. Sum the valence electrons from each
atom to get the total number of valence
electrons.
Carbon is in Group IVA (14), so it has 4
valence electrons. Each hydrogen
contributes 1 valence electron (H is in Group
IA (1)). Oxygen contributes 6 valence
electrons because it is in Group VIA (16).
Total number of valence electrons = 4 + (1 x 2) + 6
= 12
8 - 34
3. Next, bond the electrons around each atom in a
single bond first, then use double bonds as
necessary.
O
C
H
H
H
C
O
H
H
C
O
H
8 - 35
1. Write an atomic skeleton for HCN:
H C N
electrons.
Carbon is in Group IVA (14), so it has 4
valence electrons. Hydrogen contributes 1
valence electron (H is in Group IA (1)).
Nitrogen contributes 5 valence electrons
because it is in Group VA (15).
Total number of valence electrons = 4 + 1 + 5 = 10
8 - 36
12
3. Next, bond the electrons around each
atom in a single bond first, then use
double and triple bonds as necessary.
H
C
N
H C N
H C N
H C N
8 - 37
Resonance Structures
• When there are several equally valid
arrangements of bonding (i.e. Lewis Dot
structures), then the concept of
resonance helps explain why.
• Resonance
– The electron arrangement in molecules with
several equally valid Lewis Dot structures is
represented by them all, each showing a
different arrangement of the true
arrangement of electrons.
• Resonance hybrid
– Representation of the actual molecule
– A composite of the formulas drawn
8 - 38
Practice - Resonance Structures
1. Write an atomic skeleton for N2O:
N
N
O
electrons.
Nitrogen is in Group VA (15), so it has 5
valence electrons. Oxygen contributes 6
valence electrons because it is in Group VIA
(16).
Total number of valence electrons = (5 x 2) + 6
= 16
8 - 39
13
3. Next, bond the electrons around each atom in a
single bond first, then use double bonds as
necessary.
N N O
N N O
N N O
8 - 40
4. To draw resonance structures, rearrange
the electrons (and bonds) in the
structures.
N N O
N N O
8 - 41
Exceptions to the Octet Rule
•
Incomplete octets
– The central atom has less than eight electrons
around it.
– Ex. BH3
•
Expanded octets
– The central atom has greater than eight electrons
around it.
– Ex. PH5, SF6
•
Odd-numbered Lewis Dot structures
– The total number of electrons is odd.
– The central atom has an odd number of electrons
around it.
– Ex. NO, NO2, ClO2
8 - 42
14
Carbon Compounds
• Carbon has:
–
–
–
–
Four valence electrons
The ability to form four bonds
The ability to catenate, i.e. bond to itself
Very strong bonds when bonded to itself
• Carbon molecules are ubiquitous in nature.
8 - 43
Hydrocarbons
Figure 8.20
8 - 44
Hydrocarbons
• Compounds containing hydrogen and carbon
• Aliphatic hydrocarbons
– A class in which the bonds are all localized single,
double, and triple bonds
– Alkanes
• Hydrocarbons which contain only carbon-carbon
single bonds
– Alkenes
• Hydrocarbons which contain at least one carboncarbon double bond
– Alkynes
• Hydrocarbons which contain at least one carboncarbon triple bond
8 - 45
15
Hydrocarbons
• Aromatic hydrocarbons
– A class of hydrocarbons which has carbon
atoms arranged in a six-atom ring with
alternating single and double bonds
– Delocalized structures
Figure 8.22
8 - 46
Functional Groups
• Functional group
– A group that is introduced
into or substituted in a
hydrocarbon chain
– Gives the hydrocarbon its
characteristic properties
– The group has a
heteroatom, an atom other
C and H
Figure 8.23
• Typically O, S, and N
– Alcohol
• A hydroxyl group (-OH)
replaces a hydrogen atom in
the formula for a hydrocarbon
8 - 47
Table 8.4 Functional Groups in Hydrocarbons
Class
Functional
Group
Example
Formula
Alcohol
-OH
Ethyl alcohol
C2H5-OH
Ether
-O-
Diethyl ether
H5C2-O-C2H5
Aldehyde
-C-H
Acetaldehyde
Ketone
-C-
Acetone
Carboxylic
Acid
-C-OH
Acetic acid
Ester
-C-O-
Ethyl acetate
-N-
Methyl amine
Amine
H3C-NH2
8 - 48
16
Odors and Carbon Compounds
8 - 49
Shapes of Molecules
• The relative locations of electron
pairs around a central atom play a
large role in determining a
molecule’s 3-D shape.
• Negatively charged electrons repel
one another, so electron pairs in
different orbitals stay as far apart as
possible.
8 - 50
Shapes of Molecules
• Valence shell electron pair repulsion
(VSEPR) theory
– The tendency of electron pairs to
adjust the orientation of their orbitals
to maximize the distance between
them
– The bonded atoms and unshared pairs
are arranged around the central atom
as far apart as possible
• Bond angle
– A shape is characterized by a bond
angle between the central atom and
the atoms bonded to it
8 - 51
17
VSEPR Parent Structures
Table 8.5
8 - 52
VSEPR Derivative Structures
8 - 53
Steps for VSEPR Structures
1. Draw a Lewis formula.
2. Count the number of atoms bonded to
the central atom, and count unshared
pairs on the central atom.
3. Add the number of atoms and the
number of unshared electron pairs
around the central atom. The total
indicates the parent structure.
4. The molecular shape is derived from the
parent shape by considering only the
positions in the structure occupied by
bonded atoms.
8 - 54
18
Practice – VSEPR Structures
• What is the shape of the nitrite
ion (NO2-1)? What is the O-N-O
bond angle?
8 - 55
Practice Solutions – VSEPR
Structures
1. Draw the Lewis Dot Structure.
O
N
-1
O N
O
O
2. Count the number of atoms bonded to the
central atom and count unshared pairs on the
central atom.
N, which is the central atom in this case, has 2
atoms bonded to it, and 1 unshared pair on it.
8 - 56
Structures
3. If we add the number of atoms and
unshared pairs around the central
atom, we get the number 3. This
indicates that the parent structure is
trigonal planar.
B
120
A
B
B
8 - 57
19
Structures
4. The derived structure, which only
considers bonded atoms, is called bent
(or angular) and has a bond angle equal
to 118°, since we take 2°off the parent
structure’s bond angle to obtain the
derived structure’s bond angle.
N
O
118
O
8 - 58
Natural Applications of VSEPR Theory
• Molecular shapes are important in living
systems.
• Glycine molecules are typically found in proteins
or gelatins.
Figure 8.29
8 - 59
Natural Applications of VSEPR Theory
• Heme molecule
– Oxygen is carried throughout the body via red blood
cells containing heme molecules.
– Histidine, an amino acid in the heme molecule, just
fits into the space next to the oxygen molecule.
8 - 60
20
Polarity of Molecules
• Diatomic molecules
– Polarity lies along the plane of the bond
• Polyatomic molecules
– A nonpolar molecule is one that has all nonpolar
bonds or one that has polar bonds that cancel out
• Bonds that cancel out have equal polarities in
opposite directions
– This happens when:
» A central atom has no unshared electrons
» The atoms around the central atom all have
the same electronegativity
– A polar molecule is one that has polar bonds that DO
NOT cancel out
8 - 61
Polarity of Molecules
Figure 8.32
8 - 62
Practice – Polarity of
Molecules
• Predict whether C2H6, NO2,
CO2, SO2, and SO3 are polar
or nonpolar molecules.
8 - 63
21
Practice Solutions – Polarity of
Molecules
• Predict whether C2H6, NO2, CO2, SO2,
and SO3 are polar or nonpolar
molecules.
– C2H6 tetrahedral at each C nonpolar
– NO2 bent polar
– CO2 linear nonpolar
– SO2 bent polar
– SO3 trigonal planar nonpolar
8 - 64
“Like” Dissolves “Like”
• Ionic salts and
polar liquids
dissolve better in
polar liquids than in
nonpolar liquids
• Nonpolar liquids
dissolve better in
other nonpolar
liquids than in polar
liquids
Figure 8.34
8 - 65
22

Chapter 8 Chemical Bonding

Transcription

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