Transcript UNIT 3 Chemical Bonding & Intermolecular Forces CHAPTERS 6 &10

UNIT 3
CHAPTERS 6 &10
Chemical Bonding & Intermolecular Forces
Chapter 6
Chemical Bonding
Chapter 6 – Section 1: Introduction to Chemical Bonding
Chemical Bonding
• Valence electrons are the
electrons in the outer shell
(highest energy level)
of an atom.
• A chemical bond is a mutual electrical attraction
between the nuclei and valence electrons of
different atoms that binds the atoms together.
• During bonding, valence electrons are
redistributed in ways that make the atoms
more stable.
Chapter 6 – Section 1: Introduction to Chemical Bonding
The Three Major Types of Chemical Bonding
• Ionic Bonding results from the electrical
attraction between oppositely-charged ions.
• Covalent Bonding results from the sharing of
electron pairs between two atoms.
• Metallic Bonding
results from the
attraction between
metal atoms and
the surrounding
sea of electrons.
Chapter 6 – Section 1: Introduction to Chemical Bonding
Ionic or Covalent?
• Bonding is usually somewhere between ionic
and covalent, depending on the electronegativity
difference between the two atoms.
• In polar covalent bonds, the bonded atoms have
an unequal attraction for the shared electron.
0
0.3
1.7
3.3
Chapter 6 – Section 1: Introduction to Chemical Bonding
Ionic or Covalent?
Sample Problem
Use electronegativity values (in table on pg 161)
to classify bonding between sulfur, S, and the following
elements: hydrogen, H; cesium, Cs; and chlorine, Cl.
In each pair, which atom will be more negative?
Solution:
Bonding
between
sulfur and:
hydrogen
cesium
chlorine
Electroneg.
difference
2.5 – 2.1 = 0.4
2.5 – 0.7 = 1.8
3.0 – 2.5 = 0.5
Bond type
polar-covalent
ionic
polar-covalent
More
negative
atom
sulfur
sulfur
chlorine
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Molecules
• A covalent bond is formed from shared pairs of
electrons.
• A molecule is
a neutral group
of atoms held
together by
covalent bonds.
Visual Concept
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Why Do Covalent Bonds Form?
• When two atoms form a covalent bond, their
shared electrons form overlapping orbitals.
• This gives both atoms a stable noble-gas
configuration.
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
The Octet Rule
• Atoms are the most stable when
they have completely full valence
shells (like the noble gases.)
• The Octet Rule – Compounds
tend to form so that each atom
has an octet (group of eight)
electrons in its highest energy level.
• Hydrogen is an exception to the octet rule since it
can only have two electrons in its valence shell.
Visual Concept
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Electron-Dot Notation
• Electron-dot notation
is indicated by dots
placed around the
element’s symbol.
Only the valence
electrons are shown.
Inner-shell electrons
are not shown.
Visual Concept
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Electron-Dot Notation
Sample Problem
a. Write the electron-dot notation for hydrogen.
b. Write the electron-dot notation for nitrogen.
Solution:
a. Hydrogen is in group 1. It has one valence electron.
H•
a. Nitrogen is in group 15. It has 5 valence electrons.
• •
•N •
•
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Lewis Structures
• Electron-dot notations of two or more atoms
can be combined to represent molecules.
• Unpaired electrons will pair up to form a
shared pair or covalent bond.
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Lewis Structures (continued)
• The pair of dots representing
the shared pair of electrons in
a covalent bond is often replaced
by a long dash.
• An unshared pair, also called a
lone pair, is a pair of electrons
that is not involved in bonding Shared pair
(covalent bond)
and that belongs exclusively
to one atom.
•
•
Lone
pair
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
How to Draw Lewis Structures
1. Draw the electron-dot notation for each type of
atom, and count the valence electrons.
2. Put the least electronegative atom in the center
(except H.)
3. Use electron pairs to form bonds between all atoms.
4. Make sure all atoms (except H) have octets.
5. Count the total electrons in your Lewis structure.
Does it match the number you counted in step 1?
If not, introduce multiple bonds.
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Lewis Structures
Sample Problem A
Draw the Lewis structure of iodomethane, CH3I.
Solution:
Step 1 - Draw the electron-dot notation for each
type of atom, and count the valence electrons.
••
•
•
•
•• I ••
H
C
•
•
•
C
1 x 4 e- = 4 e3H
3 x 1 e- = 3 eI
1 x 7 e- = 7 e14 e- Total
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Lewis Structures
Sample Problem A (continued)
Step 2 – Put the least electronegative atom in the center
(except H).
Step 3 – Use electron pairs to form bonds between all atoms.
Step 4 – Make sure all atoms (except H) have octets.
Step 5 – Count the total electrons. Does it match your
beginning total?
14 Total e-


H
••
H C I ••
••
H
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Multiple Covalent Bonds
• In a single covalent bond, one pair of
electrons is shared between two atoms.
• A double bond is a covalent bond in
which two pairs of electrons are
shared between two atoms.
• A triple bond is a covalent bond in
which three pairs of electrons are
shared between two atoms.
• Multiple bonds are often found in
molecules containing carbon, nitrogen,
and oxygen.
Single Bond
Double Bond
Triple Bond
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Lewis Structures
Sample Problem B
Draw the Lewis structure for methanal, CH2O.
Solution:
Step 1 - Draw the electron-dot notation for each
type of atom, and count the valence electrons.
•
•
•
•
•• O ••
H
C
•
•
•
C
1 x 4 e- = 4 e2H
2 x 1 e- = 2 eO
1 x 6 e- = 6 e12 e- Total
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Lewis Structures
Sample Problem B (continued)
Step 2 – Put the least electronegative atom in the center
(except H).
Step 3 – Use electron pairs to form bonds between all atoms.
Step 4 – Make sure all atoms (except H) have octets.
Step 5 – Count the total electrons. Does it match your
beginning total? 14 Total eIf not, introduce multiple bonds (remove
••
2 lone pairs to make 1 shared pair.)
••
••
Now does it match? 12 Total e••



H
C O
H
Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds
Formation of Ionic Compounds
• Sodium and other metals easily lose electrons
to form positively-charged ions called cations.
• Chlorine and other non-metals easily gain
electrons to form negatively-charged ions
called anions.
Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds
Ionic Bonding
• Cations (+) and anions (-)
are attracted to each
other because of their
opposite electrical charges.
• An ionic bond is a bond
that forms between
oppositely-charged ions
because of their mutual
electrical attraction.
Visual Concept
Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds
Ionic Bonding and the Crystal Lattice
• In an ionic crystal, ions minimize their potential
energy by combining in an orderly arrangement
known as a crystal lattice.
• A formula unit is the smallest repeating unit of an
ionic compound.
Sodium Chloride crystal lattice (many Na and Cl atoms)
Formula Unit = NaCl
Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds
Comparing Ionic and Covalent Compounds
• Covalent compounds have relatively weak forces of
attraction between molecules, but ionic compounds
have a strong attraction between ions. This causes
some differences in their properties:
Ionic
crystals
very high melting points
hard, but brittle
Ex: NaCl, CaF2, KNO3
Covalent
molecules
low melting points
usually gas or liquid
Ex: H2O, CO2, O2
Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds
Polyatomic Ions
• A charged group of covalently bonded atoms
is known as a polyatomic ion.
• Draw a Lewis structure for a polyatomic ion
with brackets around it and the charge in the
upper right corner.
hydroxide ion, OH-
ammonium ion, NH4+
Chapter 6 – Section 4: Metallic Bonding
The Metallic Bond
• In metals, overlapping orbitals allow the outer
electrons of the atoms to roam freely
throughout the entire metal.
• These mobile electrons form
a sea of electrons around the
metal atoms, which are packed
together in a crystal lattice.
• A metallic bond results from the attraction
between metal atoms and the surrounding
sea of electrons.
Chapter 6 – Section 4: Metallic Bonding
Properties of Metals
• The characteristics of metallic bonding gives
metals their unique properties, listed below.
 electrical conductivity
 thermal (heat) conductivity
 malleability (can be
hammered into thin sheets)
 ductility (can be pulled
or extruded into wires)
 luster (shiny appearance)
Visual Concept
Chapter 6 – Section 5: Molecular Geometry
VSEPR Theory
• The abbreviation VSEPR (say it
“VES-pur”) stands for “valence-shell
electron-pair repulsion.”
• VSEPR theory – repulsion between
pairs of valence electrons around
an atom causes the electron pairs to
be oriented as far apart as possible.
• Treat double and triple bonds the same as
single bonds.
Visual Concept
Chapter 6 – Section 5: Molecular Geometry
VSEPR Theory (continued)
• VSEPR theory can also account for the geometries
of molecules with unshared electron pairs.
• VSEPR theory postulates that the lone pairs occupy
space around the central atom just like bonding
pairs, but they repel other electron pairs more
strongly than bonding pairs do.
Chapter 6 – Section 5: Molecular Geometry
VSEPR Theory (continued)
• 2 electron pairs around a
central atom will be 180o
apart, and the molecule’s
shape will be linear.
• 3 bonding pairs around a
central atom will be 120o
apart, and the molecule’s
shape will be trigonal planar.
If one of the pairs is a lone
pair, the shape will be bent.
Chapter 6 – Section 5: Molecular Geometry
VSEPR Theory (continued)
• 4 bonding pairs around a
central atom will be 109.5o
apart, and the molecule’s
shape will be tetrahedral.
If one of the pairs is a lone
pair, the shape will be
trigonal pyramidal. If two
of the pairs are lone pairs,
the shape will be bent.
• Unshared pairs repel electrons
more strongly and will result in
smaller bond angles.
Chapter 6 – Section 5: Molecular Geometry
VSEPR Theory
Sample Problem A
Use VSEPR theory to predict the molecular geometry
of water, H2O.
Solution:
Draw the Lewis Structure for H2O: Total Electrons: 8 eOctets
How many total electron pairs are
•
•
surrounding the central atom?
4
••
How many are unshared pairs?
•• ••
2
The shape is bent.

O
H
H
O

H
H
Chapter 6 – Section 5: Molecular Geometry
VSEPR Theory
Sample Problem B
Use VSEPR theory to predict the molecular geometry
of carbon dioxide, CO2.
Solution:
Total Electrons: 16 eDraw the Lewis Structure for CO2: Octets
••
••
How many total electron pairs are • • •
••
•
surrounding the central atom?
••
••

O C
••
2 (double or triple bonds count the same as single)
The shape is linear.

O
Chapter 6 – Section 5: Molecular Geometry
Molecular Polarity
• Molecular Polarity
depends on both bond
polarity and molecular
geometry.



If all bonds are
non-polar, the molecule
is always non-polar.
If bonds are polar, but there is symmetry in the
molecule so that the polarity of the bonds cancels
out, then the molecule is non-polar. (Ex: CO2, CCl4)
If bonds are polar but there is no symmetry such
that they cancel each other out, the overall
molecule is polar. (Ex: H20, CH3Cl)
Chapter 6 – Section 5: Molecular Geometry
Intermolecular Forces
• The forces of attraction between molecules
are called intermolecular forces.
• Intermolecular forces vary in strength but are
generally weaker than
any of the three types
of chemical bonds
(covalent, ionic or metallic.)
Chapter 6 – Section 5: Molecular Geometry
Intermolecular Forces (continued)
• The strongest intermolecular forces
exist between polar molecules.
• Because of their uneven
charge distribution, polar
molecules have dipoles.
• A dipole is represented by an
arrow with its head pointing
toward the negative pole and a
crossed tail at the positive pole.
Visual Concept
•• ••
H
O
H
Chapter 6 – Section 5: Molecular Geometry
Types of Intermolecular Forces
• 3 types of intermolecular forces (strongest to weakest):
1. Dipole-dipole – between 2 polar
molecules. The - side of 1 dipole
attracts the + side of another.
•
2.
3.
Hydrogen Bonding – a very strong
type of dipole-dipole force. Only exists
between atoms of H and N, O or F.
Induced dipole – between a polar
and a non-polar molecule.
London dispersion forces – instantaneous dipoles
created by the constant motion of electrons.
Visual Concept
Chapter 10
States of Matter
Chapter 10 – Section 1: The Kinetic-Molecular Theory of Matter
The Kinetic-Molecular Theory
• The kinetic-molecular theory of matter states:




Particles of matter (atoms and
molecules) are always in motion.
We measure this energy of motion
(kinetic energy) as temperature.
If temperature increases, the
particles will gain more energy
and move even faster.
Molecular motion is greatest in
gases, less in liquids, and least
in solids.
Chapter 10 – Section 1: The Kinetic-Molecular Theory of Matter
Gases
• An Ideal Gas is a hypothetical gas that
perfectly fits all the assumptions of the kineticmolecular theory.
• Many gases behave nearly
ideally if pressure is not very
high and temperature is not
very low.
• Fluidity – Gas particles glide easily past one
another. Because liquids and gases flow, they
are both referred to as fluids.
Chapter 10 – Section 1: The Kinetic-Molecular Theory of Matter
Gases (continued)
• Low Density – Gas particles are very far apart.
The density of a gas is about 1/1000 the density
of the same substance in the liquid or solid state.
• Expansion – A gas will expand to fill its container.
• Compressibility – The volume of a gas can be
greatly decreased by pushing the particles closer
together.
Chapter 10 – Section 2: Liquids
Liquids
• Surface Tension – Strong
cohesive forces at a liquid’s
surface act to decrease the
surface area to the smallest
possible size. The higher the
force of attraction between
the particles of a liquid, the
higher the surface tension.
Chapter 10 – Section 2: Liquids
Liquids (continued)
• Vaporization – A liquid or solid
changing to a gas.
Evaporation – particles escape
from the surface of a liquid and
become a gas. This occurs
because liquid particles have
different kinetic energies.
 Boiling – bubbles of vapor appear throughout a
liquid. Will not occur below a certain
temperature (the boiling point.)

• A volatile liquid is one that evaporates readily.
Chapter 10 – Section 3: Solids
Solids
• There are two main types of solids:


Crystalline Solids – Made up
of crystals. Particles are
arranged in an orderly,
geometric, repeating pattern.
Amorphous Solid – Particles
are arranged randomly.
Chapter 10 – Section 3: Solids
Solids (continued)
• Melting Point – The temperature at which a
solid becomes a liquid. At this temperature, the
kinetic energies of the particles within the solid
overcome the attractive forces holding them
together.