Chemical Bonding and Molecular Structure

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Chapter 6 Chemical Bonding and Molecular
Structure
Cartoon courtesy of NearingZero.net
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Chemical Bond
A mutual electrical attraction between the
nuclei and valence electrons of different atoms
that bind them together.
Intra-molecular forces are forces between atoms
which include 1.Covalent bonds sharing of
electrons 2.Ionic bonds - Transfer of electrons
Forces between cations and anions. 3.Metallic
bonds forces between atoms of the same metal.
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Electro-negativity

• A measure of an atoms ability to attract
  electrons in a chemical bond.
• Trend increases up and to the right on the
  periodic table.
• Values range from .7 to 4
• Fluorine has the highest value
• Differences in electro-negativity
• values can help predict bond types

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Bonding is not usually purely ionic or
covalent in character.Bonding electrons will be
more strongly attracted to atoms of higher
electro-negativity
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Predicting bond types

• 1. General rules
• Metal Nonmetal ionic
• Metal Metalloid covalent
• Nonmetal Nonmetal covalent
• Metal Metal metallic
• Use Differences in electro-negativity values
• Note- The electro-negativity ranges above are
  in reference to the Modern text book. Many
  exceptions apply and different books have
  different cut off points. These difference in
  electro-negativity will mainly help to predict
  between polar and nonpolar covalent bonds.

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1) Covalent Bonds Two types of covalent
bonds A. Polar covalent bonds A bond in which
the bonded atoms have an unequal attraction for
the shared electrons. Positive and negative poles
form on the molecule. B. Nonpolar covalent
bonds A bond in which the bonding electrons are
shared equally by the bonded atoms resulting in a
balanced distribution of electrical charge. No
poles are formed on the molecule.
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• Sample Problem A
• Use electro-negativity values listed in Figure 20
  from the previous chapter in your book, on page
  161, and Figure 2 in your book, on page 176, to
  classify bonding between sulfur and the following
  elements
• hydrogen
• cesium
• chlorine
• In each pair, which atom will be more negative?

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Sample Problem A Solution The electro-negativity
of sulfur is 2.5. The electro-negativities of
hydrogen, cesium, and chlorine are 2.1, 0.7, and
3.0, respectively. In each pair, the atom with
the larger electro-negativity will be the
more-negative atom.
Bonding between More
sulfur and eneg Eneg
difference Bond type element hydrogen 2.5 2.1
0.4 polar-covalent sulfur cesium 2.5 0.7
1.8 ionic sulfur chlorine 3.0 2.5
0.5 polar-covalent chlorine
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Bond Type Prediction
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Ionic Bonds

• Definition composed of positive and negative
  ions that are combined so that the number of
  positive and negative charges are equal.
• A formula unit is the smallest unit of atoms from
  which an ionic compounds formula can be
  established.
• Most ionic compounds exist as crystalline solids.
• A crystal of any ionic compound is a
  three-dimensional network of positive and
  negative ions mutually attracted to each other.

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Ionic Compounds

• The chemical formula of an ionic compound
  represents not molecules, but the simplest ratio
  of the compounds ions.

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The Octet Rule Ionic Compounds
Ionic compounds tend to form so that each atom
has an octet of electrons by gaining or losing in
its highest occupied energy level.
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Ionic BondingThe Formation of Sodium Chloride

• Sodium has 1 valence electron

• Chlorine has 7 valence electrons

• An electron transferred gives
• each an octet

Na 1s22s22p63s1
Cl 1s22s22p63s23p5
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Ionic BondingThe Formation of Sodium Chloride
This transfer forms ions, each with an octet
Na 1s22s22p6
Cl- 1s22s22p63s23p6
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Ionic BondingThe Formation of Sodium Chloride
The resulting ions come together due to
electrostatic attraction (opposites attract)
Cl-
Na
The net charge on the compound must equal zero
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Examples of Ionic compounds
All salts, which are composed of metals bonded
to nonmetals, are ionic compounds and form ionic
crystals.
Examples
Na1O-2
Ca2O2-
Mg2Cl-1
Al3O-2
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Properties of Ionic Compounds
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Sodium Chloride Crystal Lattice
Ionic compounds form solids at ordinary
temperatures.
Ionic compounds organize in a characteristic
crystal lattice of alternating positive and
negative ions.
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Representation of Components in an Ionic Solid
Lattice A 3-dimensional system of points
designating the centers of components (atoms,
ions, or molecules) that make up the substance.
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Covalent Molecular Compounds

• Definition results from the sharing of electron
  pairs between two atoms.
• A molecule is a neutral group of atoms that are
  held together by covalent bonds.
• A chemical compound whose simplest units are
  molecules is called a molecular compound.

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• A chemical formula indicates the relative numbers
  of atoms of each kind in a chemical compound by
  using atomic symbols and numerical subscripts.
• H2O
• A molecular formula shows the types and numbers
  of atoms combined in a single molecule of a
  molecular compound. The composition of a compound
  is given by its chemical formula.
• H2O

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Diatomic Molecules A diatomic molecule is a
molecule that contains only 2 atoms. The
diatomic 7 are I2, Br2, Cl2, F2, O2, N2,
H2! Think about our song!
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Covalent Bond

• Most atoms have lower potential energy when they
  are bonded to other atoms than they have as they
  are independent particles.
• The figure below shows potential energy changes
  during the formation of a hydrogen-hydrogen bond.
  The lowest potential energy usually creates the
  most stability

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Characteristics of the Covalent Bond

• The average distance between two bonded atoms at
  their lowest potential energy is called bond
  length.
• In forming a covalent bond, the hydrogen atoms
  release energy. The same amount of energy must be
  added to separate the bonded atoms.
• Bond energy is the energy required to break a
  chemical bond and form neutral isolated atoms.
• The stronger the bond, the shorter the bond
  length (distance) and the higher the bond energy
  value.
• An inverse proportion exist between bond energy
  and length.

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Bond Length and Stability
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Ionic Vs. Covalent Bonding

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A Comparison of Ionic and Molecular Compounds

• The force that holds ions together in an ionic
  compound is a very strong electrostatic
  attraction.
• In contrast, the forces of attraction between
  molecules of a covalent compound are much weaker.
• This difference in the strength of attraction
  between the basic units of molecular and ionic
  compounds gives rise to different properties
  between the two types of compounds.

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A Comparison of Ionic and Molecular Compounds

• Molecular compounds have relatively weak forces
  between individual molecules.
• They melt at low temperatures.
• The strong attraction between ions in an ionic
  compound gives ionic compounds some
  characteristic properties, listed below.
• very high melting points
• hard but brittle
• not electrical conductors in the solid state,
  because the ions cannot move good in molten or
  aqueous states.

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The Octet Rule for Covalent Compounds
Covalent compounds tend to form so that each
atom, by sharing electrons, has an octet (8) of
electrons in its highest occupied energy level.
Diatomic Fluorine
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Hydrogen Chloride by the Octet Rule
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Electron Dot Notation Dots represent valence
electrons. Valence electrons are found in the s
and p of the highest main energy level.
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Dot Notation
Sample Problem B a. Write the electron-dot
notation for hydrogen. b. Write the electron-dot
notation for nitrogen.
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Sample Problem B Solution a. A hydrogen atom
has only one occupied energy level, the n 1
level, which contains a single electron. Group IA

b. The group notation for nitrogens
family of elements is ns2np3. Group VA.
Nitrogen has five valence electrons.
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Formation of Water by the Octet Rule
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Comments About the Octet Rule

•  2nd row elements C, N, O, F observe the octet
  rule.
• Many exceptions apply
• H defines this rule and only needs two electrons.
•  2nd row elements B and Be often have fewer than
  8 electrons around themselves - they are very
  reactive. Be can have 4,6 8 and B can have 6 or
  8.
•  3rd row and heavier elements CAN exceed the
  octet rule using empty valence d orbitals.
•  When writing Lewis structures, satisfy octets
  first, then place electrons around elements
  having available d orbitals.

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Lewis Dot Structures

• Electron-dot notation can also be used to
  represent molecules.

• The pair of dots between the two symbols
  represents the shared electron pair of the
  hydrogen-hydrogen covalent bond.
• For a molecule of fluorine, F2, the electron-dot
  notations of two fluorine atoms are combined.

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Lewis Dot Structures

• The pair of dots between the two symbols
  represents the shared pair of a covalent bond.

• In addition, each fluorine atom is surrounded by
  three pairs of electrons that are not shared in
  bonds.

• 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.

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Lewis Structures

•  Shows how valence electrons are arranged among
  atoms in a molecule.
•  Reflects central idea that stability of a
  compound relates to noble gas electron
  configuration.

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Lewis Structures

• The pair of dots representing a shared pair of
  electrons in a covalent bond is often replaced by
  a long dash.
• example

• A structural formula indicates the kind, number,
  and arrangement, and bonds but not the unshared
  pairs of the atoms in a molecule.
• example FF HCl

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Single Covalent Bonds
A single covalent bond, or single bond, is a
covalent bond in which one pair of electrons is
shared between two atoms. Example H-H
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Multiple Covalent BondsDouble bonds
Two pairs of shared electrons
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Multiple Covalent BondsTriple bonds
Three pairs of shared electrons
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Drawing Lewis Structures with many atoms

• Determine the type and number of atoms in the
  molecule.
• Determine the total number of valence electrons
  to be used.
• Form a skeletal arrangement. Symmetrical is more
  likely.
• C or the least electronegative element is usually
  the central atom.
• H is never the central atom
• Place unshared electron pairs so that each atom
  is surrounded by eight electrons.
• Recount the total valence electrons.

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Sample Problem C Draw the Lewis structure of CH3I.
1. Determine the type and number of atoms in the
molecule. The formula shows one carbon atom, one
iodine atom, and three hydrogen atoms. 2. Write
the electron-dot notation for each type of atom
in the molecule and determine the total number of
valence electrons available in the atoms to be
combined. Carbon is from Group 14 (IVA) and
has four valence electrons. Iodine is from Group
17 (VIIA) and has seven valence electrons.
Hydrogen has one valence electron. There are
three H atoms.
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• Determine the total number of valence electrons
  available in the atoms to be combined.
• C 1 x 4 4
• I 1 x 7 7
• H 3 x 1 3
• 14 valence electrons
  available

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• If carbon is present, it is the central atom.
  Otherwise, the least-electronegative atom is
  central. Hydrogen, is never central.

5. Add unshared pairs of electrons to each
nonmetal atom (except hydrogen) such that each is
surrounded by eight electrons.
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Sample Problem D Draw the Lewis structure for
methanol, CH2O, which is also known as
formaldehyde.
1. Determine the number of atoms of each element
present in the molecule. The formula shows one
carbon atom, two hydrogen atoms, and one oxygen
atom. 2. Write the electron-dot notation for each
type of atom and determine the total number of
valence electrons available in the atoms to be
combined Carbon is from Group 14 (IVA) and has
four valence electrons. Oxygen, which is in Group
16 (VIA) and has six valence electrons. Hydrogen
has only one valence electron.
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• Determine the total number of valence electrons
  available in the atoms to be combined.
• C - 1 x 4 4
• H - 2 x 1 2
• O - 1 x 6 6
• 12 valence electrons
  available

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• Arrange the atoms to form a skeleton structure
  for the molecule. Connect the atoms by
  electron-pair bonds.

5. Add unshared pairs of electrons to each
nonmetal atom (except hydrogen) such that each is
surrounded by eight electrons.
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6a.Count the electrons in the Lewis structure to
be sure that the number of valence electrons used
equals the number available. The structure has
14 electrons. The structure has two valence
electrons too many. 6b.Subtract one or more
lone pairs until the total number of valence
electrons is correct. Move one or more lone
electron pairs to existing bonds until the outer
shells of all atoms are completely filled.
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Subtract the lone pair of electrons from the
carbon atom. Move one lone pair of electrons from
the oxygen to the bond between carbon and oxygen
to form a double bond.
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Completing a Lewis Structure of CH3Cl

• Make carbon the central atom

• Add up available valence electrons

• C 4, H (3)(1), Cl 7 Total 14

•   Join peripheral atoms
• to the central atom
• with electron pairs.

H
..
..
..
H
..
..
C

•   Complete octets on
• atoms other than
• hydrogen with remaining
• electrons

Cl
..
..
H
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Lewis Structures Practice Drawing and Labeling

• Draw Lewis Structures

• Terms to mention and label

• Polar covalent bond
• Non polar covalent bond
• Lone pairs
• Multiple bonds
• Double covalent bond
• Triple covalent bond
• Single covalent bond

• F and F
• H and Cl
• Br and Br
• O and O
• N and N
• CH3I
• NH3
• H2
• SiF4
• H2S
• H2O

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Resonance

•  Occurs when more than one valid Lewis structure
  can be written for a particular molecule.

•  These are resonance structures.
• The actual structure is an average of
• the resonance structures.
• In a benzene ring the double
• and single bonds resonate.

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Draw the resonant structures for the polyatomic
ion CO3-2
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VSEPR Model
Molecular Geometry for covalent
compounds (Valence Shell Electron Pair
Repulsion) VSEPR THEORY

•  The structure around a given atom is determined
  principally by minimizing electron pair
  repulsions in covalent compounds.

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Predicting a VSEPR Structure

•   Draw Lewis structure.
•   Put pairs as far apart as possible.
•   Determine positions of atoms from the way
  electron pairs are shared.
•   Determine the name of molecular structure from
  positions of the atoms.

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VSEPR Theory

• Unshared electron pairs repel other electron
  pairs more strongly than bonding pairs do. This
  is why the bond angles in ammonia and water are
  somewhat less than the 109.5 bond angles of a
  perfectly tetrahedral molecule.

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Determining Bond Polarity for VSEPR Structures
To determine bond polarity, simply take the
electro-negativity difference (using
electro-negativity table) in bonding atoms and
determine into what range that this difference
belongs Polar covalent or nonpolar covalent.
Ex H and Cl Dipole


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Determining Molecule Polarity for VSEPR
Structures
To determine molecule polarity, look at the
molecule to determine if there are regions of
uneven sharing of electrons. If so, then the
molecule is polar. If all electrons are being
shared equally all over the molecule, the
molecule is nonpolar.
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Polarity
A molecule, such as HF, that has a positively
charged end and a negatively charged end is said
to be a polar molecule, or to have a dipole
moment. The bond on HF is also polar.
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Rules for Predicting Molecule Polarity

• For diatomic and 2 atom molecules, the bond
  polarity indicates the molecule polarity.
• For molecules with 3 or more atoms, if the
  dipoles of the individual bonds cancel each
  other, the molecule is nonpolar.
• Linear, trigonal and tetrahedral are nonpolar, if
  the same terminal element is bonded.
• Bent and pyramidal with lone pairs are always
  polar.
• Multiple bonds have no effect on geometric shape

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VSEPR and Molecular Geometry
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VSEPR and Molecular Geometry
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The Octet Rule Ionic Compounds
Ionic compounds tend to form so that each atom,
by electrons, has an octet of electrons gaining
or losing in its highest occupied energy level.
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Polyatomic Ions

• Certain atoms bond covalently with each other to
  form a group of atoms that has both molecular and
  ionic characteristics.
• A charged group of covalently bonded atoms is
  known as a polyatomic ion.
• Like other ions, polyatomic ions have a charge
  that results from either a shortage or excess of
  electrons. Two polyatomic ions are ionically
  bonded.

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Polyatomic Ions
An example of a polyatomic ion is the ammonium
ion NH41. It is sometimes written as NH41
to show that the group of atoms as a whole has
a charge of 1.

• The charge of the ammonium ion is determined as
  follows
• Draw the bonding in ammonium fluoride.
• ?????
• The seven protons in the nitrogen atom plus the
  four protons in the four hydrogen atoms give the
  ammonium ion a total positive charge of 11.

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Polyatomic Ions

• Some examples of Lewis structures of polyatomic
  ions are shown below.

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3) Metallic Bonding

• The chemical bonding that results from the
  attraction between metal atoms and the
  surrounding sea of electrons.
• Only positive ions exist in metallic bonds
• Vacant p and d orbitals in metal's outer energy
  levels overlap, and allow outer electrons to roam
  freely throughout the entire metal. They are
  said to be delocalized which means that they do
  not belong to any one atom but move freely about
  the empty atomic orbitals.

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Properties of Metals

• Metals are good conductors of heat and
  electricity
• Metals are malleable
• Metals are ductile
• Metals have high tensile strength
• Metals have high luster

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Packing in Metals
Model Packing uniform, hard spheres to best use
available space. This is called closest packing.
Each atom has 12 nearest neighbors.
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Metal Alloys

• Substitutional Alloy some metal atoms replaced
  by others of similar size.

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Metal Alloys

• Interstitial Alloy Interstices (holes) in
  closest packed metal structure are occupied by
  small atoms.

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Properties of Substances with Metallic, Ionic,
and Covalent Bonds
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Polarity To determine if a molecule is polar or
non-polar, look at each dipole. A dipole is
created by equal but opposite charges that are
separated by a short distance. The direction of
a dipole is from the dipoles positive pole to
its negative pole.
Look at HCl. Because chlorine is more
eletronegative, the dipole looks like H Cl
The overall molecule is considered polar because
there is a positive end and a negative end.
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Determine if the following molecules are polar or
non-polar. Be sure and look at individual
dipoles.
Ex. 1 H2O
The overall molecule is polar!

• Lets look at some other examples. Determine if
  the following are polar or non-polar.
• LiCl 2. CH4

3. NH3 4. CO2
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Intermolecular Forces

• The forces of attraction between molecules are
  known as intermolecular forces. Intermolecular
  forces are usually represented by dotted lines
  between molecules. Includes 3 types

• Hydrogen bonds (strongest dipoles)
• occurs when H is attracted to N,O, F
• 2. Dipole-Dipole
• 3. Dispersion
• All inter-molecular forces are
• always weaker than any of the
• intra-molecular forces.

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1. Hydrogen Bonding The intermolecular force in
which a hydrogen atom that is bonded to a highly
electronegative atom is attracted to an unshared
pair of electrons of an electronegative atom in a
nearby molecule is known as hydrogen bonding.
Hydrogen bonds -The strongest intermolecular
forces that exist between polar covalent
molecules. Ex HF, H2O
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Hydrogen Bonding in Water
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Hydrogen Bonding

• Some hydrogen-containing compounds have unusually
  high boiling points. This is explained by a
  particularly strong type of dipole-dipole force.
• In compounds containing HF, HO, or HN bonds,
  the large electronegativity differences between
  hydrogen atoms and the atoms they are bonded to
  make their bonds highly polar.
• This gives the hydrogen atom a positive charge
  that is almost half as large as that of a bare
  proton.

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2. Dipole-Dipole forces The forces of attraction
between polar molecules are known as
dipole-dipole.
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Dipole-Dipole

• The negative region in one polar molecule
  attracts the positive region in adjacent
  molecules. So the molecules all attract each
  other from opposite sides.
• Such forces of attraction between polar molecules
  are known as dipole-dipole forces.
• Dipole-dipole forces act at short range, only
  between nearby molecules.
• Dipole-dipole forces explain, for example the
  difference between the boiling points of iodine
  chloride, ICl (97C), and bromine, BrBr (59C).

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Comparing Dipole-Dipole Forces
Molecular Geometry
Chapter 6
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3. London Dispersion Forces The intermolecular
attractions resulting from the constant motion of
electrons and the creation of instantaneous
dipoles are called London dispersion forces.

• London Dispersion Forces - Even noble gas atoms
  and nonpolar molecules can experience weak
  intermolecular attraction.
• In any atom or moleculepolar or nonpolarthe
  electrons are in continuous motion.
• As a result, at any instant the electron
  distribution may be uneven. A momentary uneven
  charge can create a positive pole at one end of
  an atom of molecule and a negative pole at the
  other.

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London Dispersion Forces
This temporary dipole can then induce a dipole in
an adjacent atom or molecule. The two are held
together for an instant by the weak attraction
between temporary dipoles. Fritz London first
proposed their existence in 1930.
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Intermolecular forces and boiling points

• The boiling point of a liquid is a good measure
  of the intermolecular forces between its
  molecules the higher the boiling point, the
  stronger the forces between the molecules.
• Intermolecular forces vary in strength but are
  generally weaker than bonds between atoms within
  molecules, ions in ionic compounds, or metal
  atoms in solid metals.
• Boiling points for ionic compounds and metals
  tend to be much higher than those for molecular
  substances forces between molecules are weaker
  than those between metal atoms or ionic ions.