UNIT 3 Chapters 6

1
UNIT 3Chapters 6 10

• Chemical Bonding Intermolecular Forces

2
Chapter 6

• Chemical Bonding

3
Chemical Bonding
Chapter 6 Section 1 Introduction to 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.

4
The Three Major Types of Chemical Bonding
Chapter 6 Section 1 Introduction to 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.

5
Ionic or Covalent?
Chapter 6 Section 1 Introduction to Chemical
Bonding

• 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.

0.3
0
1.7
3.3
6
Ionic or Covalent?Sample Problem
Chapter 6 Section 1 Introduction to Chemical
Bonding

• 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 Morebetween Electroneg.
negative sulfur and difference Bond type atom
2.5 2.1 0.4
polar-covalent
sulfur
hydrogen cesium chlorine
2.5 0.7 1.8
ionic
sulfur
3.0 2.5 0.5
polar-covalent
chlorine
7
Molecules
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

• A covalent bond is formed from shared pairs of
  electrons.
• A molecule is a neutral group of atoms held
  together by covalent bonds.

8
Why Do Covalent Bonds Form?
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

• When two atoms form a covalent bond, their shared
  electrons form overlapping orbitals.
• This gives both atoms a stable noble-gas
  configuration.

9
The Octet Rule
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

• Atoms are the most stable whenthey have
  completely full valenceshells (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.

10
Electron-Dot Notation
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

• Electron-dot notation is indicated by dots
  placed around the elements symbol. Only the
  valence electrons are shown.Inner-shell
  electrons are not shown.

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Electron-Dot NotationSample Problem
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

• a. Write the electron-dot notation for hydrogen.
• b. Write the electron-dot notation for nitrogen.
• Solution
• Hydrogen is in group 1. It has one valence
  electron.
• Nitrogen is in group 15. It has 5 valence
  electrons.

H


N




12
Lewis Structures
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

• 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.

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Lewis Structures (continued)
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

• 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 and that belongs exclusively to one
  atom.

Shared pair (covalent bond)
Lone pair
14
How to Draw Lewis Structures
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

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.

15
Lewis StructuresSample Problem A
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

• 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.

C 1 x 4 e- 4 e-
3H 3 x 1 e- 3 e-
I 1 x 7 e- 7 e-
14 e- Total
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Lewis StructuresSample Problem A (continued)
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

• 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?

?
H
?
14 Total e-


H
C
I




H
17
Multiple Covalent Bonds
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

• In a single covalent bond, one pair ofelectrons
  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
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Lewis StructuresSample Problem B
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

• 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.

C 1 x 4 e- 4 e-
2H 2 x 1 e- 2 e-
O 1 x 6 e- 6 e-
12 e- Total
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Lewis StructuresSample Problem B (continued)
Chapter 6 Section 2 Covalent Bonding and
Molecular Compounds

• 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?
• If not, introduce multiple bonds (remove 2 lone
  pairs to make 1 shared pair.)
• Now does it match?

?
?
14 Total e-
H


C
O


?


12 Total e-


H
20
Formation of Ionic Compounds
Chapter 6 Section 3 Ionic Bonding and 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.

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Ionic Bonding
Chapter 6 Section 3 Ionic Bonding and Ionic
Compounds

• Cations () and anions (-)are attracted to each
  other because of their opposite electrical
  charges.
• An ionic bond is a bondthat forms
  betweenoppositely-charged ionsbecause of their
  mutualelectrical attraction.

22
Ionic Bonding and the Crystal Lattice
Chapter 6 Section 3 Ionic Bonding and Ionic
Compounds

• 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
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Comparing Ionic and Covalent Compounds
Chapter 6 Section 3 Ionic Bonding and Ionic
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
Covalent
molecules
crystals
very high melting points
low melting points
hard, but brittle
usually gas or liquid
Ex NaCl, CaF2, KNO3
Ex H2O, CO2, O2
24
Polyatomic Ions
Chapter 6 Section 3 Ionic Bonding and Ionic
Compounds

• 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
25
The Metallic Bond
Chapter 6 Section 4 Metallic Bonding

• 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.

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Properties of Metals
Chapter 6 Section 4 Metallic Bonding

• 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)

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VSEPR Theory
Chapter 6 Section 5 Molecular Geometry

• The abbreviation VSEPR (say it VES-pur) stands
  for valence-shell electron-pair repulsion.
• VSEPR theory repulsion between pairs of
  valence electrons aroundan atom causes the
  electron pairs tobe oriented as far apart as
  possible.
• Treat double and triple bonds the same as single
  bonds.

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VSEPR Theory (continued)
Chapter 6 Section 5 Molecular Geometry

• 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.

29
VSEPR Theory (continued)
Chapter 6 Section 5 Molecular Geometry

• 2 electron pairs around acentral atom will be
  180oapart, and the moleculesshape will be
  linear.
• 3 bonding pairs around acentral atom will be
  120oapart, and the moleculesshape will be
  trigonal planar.If one of the pairs is a
  lonepair, the shape will be bent.

30
VSEPR Theory (continued)
Chapter 6 Section 5 Molecular Geometry

• 4 bonding pairs around acentral atom will be
  109.5oapart, and the moleculesshape will be
  tetrahedral.If one of the pairs is a lonepair,
  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.

31
VSEPR TheorySample Problem A
Chapter 6 Section 5 Molecular Geometry

• Use VSEPR theory to predict the molecular
  geometry of water, H2O.
• Solution
• Draw the Lewis Structure for H2O
• How many total electron pairs aresurrounding the
  central atom?
• How many are unshared pairs?
• The shape is bent.

?
Total Electrons 8 e-
?
Octets


H
H
O
4






O
2
H
H
32
VSEPR TheorySample Problem B
Chapter 6 Section 5 Molecular Geometry

• Use VSEPR theory to predict the molecular
  geometry of carbon dioxide, CO2.
• Solution
• Draw the Lewis Structure for CO2
• How many total electron pairs aresurrounding the
  central atom?
• The shape is linear.

?
Total Electrons 16 e-
?
Octets






C
O
O










2 (double or triple bonds count the same as
single)
33
Molecular Polarity
Chapter 6 Section 5 Molecular Geometry

• Molecular Polarity depends on both bondpolarity
  and molecular geometry.
• If all bonds are non-polar, the moleculeis
  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)

34
Intermolecular Forces
Chapter 6 Section 5 Molecular Geometry

• 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.)

35
Intermolecular Forces (continued)
Chapter 6 Section 5 Molecular Geometry

• 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.

36
Types of Intermolecular Forces
Chapter 6 Section 5 Molecular Geometry

• 3 types of intermolecular forces (strongest to
  weakest)
• Dipole-dipole between 2 polar molecules. The -
  side of 1 dipole attracts the side of another.
• Hydrogen Bonding a very strong type of
  dipole-dipole force. Only existsbetween 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.

37
Chapter 10

• States of Matter

38
The Kinetic-Molecular Theory
Chapter 10 Section 1 The Kinetic-Molecular
Theory of Matter

• 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, theparticles will gain
  more energy and move even faster.
• Molecular motion is greatest in gases, less in
  liquids, and least in solids.

39
Gases
Chapter 10 Section 1 The Kinetic-Molecular
Theory of Matter

• An Ideal Gas is a hypothetical gas that perfectly
  fits all the assumptions of the kinetic-molecular
  theory.
• Many gases behave nearlyideally if pressure is
  not veryhigh 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.

40
Gases (continued)
Chapter 10 Section 1 The Kinetic-Molecular
Theory of Matter

• 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.

41
Liquids
Chapter 10 Section 2 Liquids

• Surface Tension Strong cohesive forces at a
  liquids 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.

42
Liquids (continued)
Chapter 10 Section 2 Liquids

• Vaporization A liquid or solidchanging to a
  gas.
• Evaporation particles escape from the surface
  of a liquid andbecome a gas. This occurs
  because liquid particles havedifferent 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.

43
Solids
Chapter 10 Section 3 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.

44
Solids (continued)
Chapter 10 Section 3 Solids

• 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.