Chemical Bonding

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Chemical Bonding
UNIT 4 Chapters 15 16
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Ionic Bonding

• The bond in ionic compounds
• (two ions)
• Held together tightly
• High melting points

3
Compounds are formed from chemically bound atoms
or ions
Substances become more stable through chemical
bonding, where 2 or more atoms are joined
together by a simultaneous attraction.
4

• Valence electrons are electrons in the highest
  occupied energy level of an atom ( the last
  shell).
• Bonding involves only electrons

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Bonding involves only the valence electrons.
Na
Cl
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Bonding Types

1. Ionic Bonding
2. Covalent Bonding
3. Metallic Bonding

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Ionic Bonds occur when the more electronegative
element steals the electron pair away from the
other atom.
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The atom that has stolen the electron pair
becomes a negative ion (anion) while the victim
becomes a positive ion (cation). The two atoms
are held together by their opposite charges.
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Can you predict which atoms will gain electrons
and which will loose electrons by looking at the
trend in electronegativity?
Increase in Electronegativity
Increase in Electronegativity
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When you consider that for an ionic bond to form
there must be a great deal of difference in
electronegativity between the atoms, can you
predict what two types of atoms allow this to
occur?
Non- Metals
Metals
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Ionic Properties
Why do most ionic compounds have similar
properties?
We can hypothesis that it is due to the bonds
formed between the ions, holding them firmly in
a rigid structure
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Forming ions

• Na and Cl which one will lose electrons which one
  will gain electrons
• Write out Tins electron configuration what will
  it do??

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Lattice energy

• The change in energy that takes place when
  separated gaseous ions are packed together to
  form an ionic solid

14
The anions and cations in an ionic compound are
locked in a regular neutrally charged structure,
held by the balance of attractive bonds and
electrical repulsion.
15

• The component ions in such crystals are arranged
  in repeating three-dimensional
• (3-D) patterns.

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Alkali metals combine with halogens in 11
ratios since alkali metals need to lose 1 e1-
and halogens need to gain 1e1-.
Alkaline earth metals combine with halogens in
12 ratios since alkaline earth metals need to
lose 2 e1- and halogens need to gain 1e1-.
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Lewis Structures can be used to illustrate the
formation of ionic bonds.

Write an equation with electron dot diagrams to
illustrate the formation of aluminum chloride.
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Lewis Structures

• Duet Rule applies to H and He and states these
  two atoms are stable with 2 electrons in their
  outer shell
• Octet Rule elements are most stable with 8
  electrons in their outer shell

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Properties of Ionic Compounds

1. Most are crystalline in structure
2. High melting/boiling points
3. Electrically neutral
4. Can conduct electricity when melted or in aqueous
   solution
5. Hard/ Brittle

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LEWIS DOT STRUCTURES Elements Board
Practice Elements 1-20
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Covalent Bonding
Br Br Br Br
O O O O
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Covalent Bonding

• Electrons are shared by nuclei
• Polar covalent bonds unequal sharing of
  electrons

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Types of Bonds

• 1) Single bond 1 pair of e- are shared
• - lowest in energy
• -longest bond length
• 2) Double bond 2 pairs of e- are shared
• 3) Triple bond- 3 pairs of e- are shared

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The most common chemical bond results when the
nuclei of 2 atoms are attracted to a pair of
shared electrons. If the sharing is equal,
because the atoms are the same, this is called
COVALENT BONDING
H H
H H
Electron pair
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• One atom becomes slightly positive the other
  slightly negative

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The force of attraction of an elements nucleus
for electrons is called electronegativity (En).
Atoms of different elements have different
electronegativities. The higher the En, the
stronger the attraction for electron pairs.
Difference in En?
HF
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The bonding electrons are on the average closer
to the fluorine than to the hydrogen atom.
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The movement of the negatively charged electrons
away from hydrogen toward fluorine, due to a
difference in electronegativity, builds up a
partial negative charge on the fluorine and a
partial positive charge on the hydrogen.
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This is not a complete transfer of an electron
from hydrogen to fluorine it is merely a
drifting of electrons toward fluorine.
F
H
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F
H
When a charge separation of this type is
present, the molecule possesses an electric
dipole, and the bond is called a POLAR COVALENT
BOND , or simply a POLAR BOND.
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F
H
Polar covalent bond (polar bond) ? covalent bond
joins two atoms of different elements and the
bonding electrons are shared unequally
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Resonance

• occurs when more than one valid Lewis structure
  can be written for a particular molecule
• Ex. NO3-1

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Exceptions to the Octet Rule

1. B and Be usually have less than 8 electrons
2. Elements in the 3rd energy level and above can
   have more than 8 electrons in their outer shell

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Non- Polar bond

• Non-polar covalent bond ? bonding electrons are
  shared equally

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Properties of Covalent Bonds

1. Soft and squishy
2. Low boiling/melting points
3. Tend to be more flammable
4. Do not conduct electricity
5. Usually non-soluble in water

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Metallic Bonding

• Electrostatic attraction force between the cation
  and free electrons.

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• Any successful bonding model for metals must
  account for the typical physical properties of
  metals malleability, ductility, and efficient
  and uniform conduction of heat and electricity in
  all directions.
• Most metals are durable and have high melting
  points.
• These facts indicate that the bonding in most
  metals are strong and nondirectional.

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• Metal atoms are arranged in very compact and
  orderly patterns.
• Body-centered cubic
• Face-centered cubic
• Hexagonal close-packed

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Properties of Metallic Compounds

1. Can conduct electricity (free electrons)
2. Malleable (put into shape)
3. Ductile ( made into wires)
4. Good conductors of heat and electricity
5. Metals are usually shiny

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Lewis Structures for Molecular Compounds
N N
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Bonding capacity is the number of covalent bonds
(shared electron pairs) that an atom can form.
Covalent molecules often consist of atoms of
different elements, with different bonding
capacities.
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How do we decide on their structural arrangement
, when we draw structural formulas?
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STEP 1place the single atom in the
center and other atoms around it evenly
spaced
H
H
H
C
H
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STEP 2 count the total of valence
e- for all atoms involved in the bonding
Carbon 1 carbon with 4 valence
electrons (1x4) 4
CH4 44 8
Hydrogen 4 hydrogen with 1
valence electrons (4x1) 4
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STEP 3 place the electrons in pairs
between the central atom and each
non-central atom
CH4 44 8
H
H
C
H
H
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STEP 4 place the remaining electrons
around the non- central atom until each
has 8 electrons (H atoms have)
only 2e-
Ex AsBr3
5 (7x3) 26
As
Br
Br
Br
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STEP 5 if electrons remain they are
placed in pairs around the central atom
Ex AsBr3
5 (7x3) 26
As
Br
Br
Br
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STEP 6 if the central atom is in group
14, 15, 16, 17 or 18, the octet rule
must be satisfied by moving electron
pairs from non-central atoms,
creating multiple bonds.
Ex SO2 6 (6x2) 18
O
O

S
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• Central B(6e-) and Be(4e-) will have
• less than 8 electrons
• 2. If the central atom is in energy
• level 3 or more it may have more
• than 8 electrons around it (these
• energy levels can have 18e-)

Cl
SCl4 6(7x4) 34
S
Cl
Cl
Cl
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NOTE central atom(s) tend to have the
highest bonding capacities or/and the lowest En.
Draw Lewis Structures for the following H2O,
NF3, Cl2, SnCl2, PCl5, SO3, BeCl2, C2H6, C2H2,
ClF3, CHCl3, ICl, O2, N2, SF6, CO2, BF3, C2H4,
O3, IF7
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ARRANGEMENT OF ATOMS IN MOLECULES

• Compounds are arranged in many different shapes

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VSEPR THEORY

• The VSEPR Theory states that because electron
  pairs repel, molecular shape adjusts so the
  valence-electron pairs are as far apart as
  possible.

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VSEPR model

• Valence shell electron pair repulsion
• Used to predict the geometry of molecules
• The structure will minimize electron pair
  repulsions

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TYPES OF MOLECULAR SHAPES
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LINEAR

• LINEAR the two bonding pairs arrange themselves
  at 180. They are connected in a straight line.
  Ex. CO2, BeF2, HCN, CS2
• Groups 2
• Pairs 0

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Tetrahedral

• A molecule that has a tetrahedral shape has all
  four pairs of electrons bonded
• Group 4
• Pair 0
• The four electron pairs repel each other forming
  an angle of 109.5

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Bent

• Molecules with a bent shape have four pairs of
  electrons, but only two pairs are bonding pairs
  (two are lone pairs). Ex H2O, SO2
• The bond angle is 109.5
• Group 2
• Pairs 2

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Trigonal Planar

• A molecule with trigonal planar shape has three
  bonds all of which lie in the same plane
• Ex. Boron trifluoride
• The bond angle are 120 .
• Group 3
• Pairs 0

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Trigonal Pyramidal

• A molecule with trigonal pyramidal shape has four
  pairs of electrons all repelling each other.
• Groups 3
• Pairs 1
• Ex. ammonia

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INTERMOLECULAR FORCES

• Intermolecular forces play a key role in
  determining the physical and chemical properties
  of covalent compounds.

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Van Der Waals

• Van Der Waals consists of 2 possible types of
  forces
• London Dispersion Forces
• Dipole-Dipole Forces

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London Dispersion Forces

• This is the only type of force present in
  non-polar covalent molecules.
• It is the weakest of the intermolecular
  interactions caused by the motion of the
  electrons.
• The strength of dispersion forces generally
  increases as the number of electrons in a
  molecules increases.
• Ex. Halogen diatomic molecules.

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Dipole-Dipole Forces

• - This occurs when polar covalent bonds are
  attracted to one another.
• - Electrostatic attractions occur between
  oppositely charged regions.
  (partially () and partially ()).
• - Dipole interactions are similar to but much
  weaker than ionic bonds.

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Dipolar molecules

• Have a center of positive charge and a center of
  negative charge.
• aka dipole moment
• Ex. HF

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(No Transcript)
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Dipole moment in NH3
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Dipole cancels out in CO2
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Hydrogen Bonding

• This is found in polar covalent molecules that
  have hydrogen that is bonded to a very
  electrostatic element (N2, F2, O2)
• Hydrogen bonds are the strongest of the
  intermolecular forces.
• Hydrogen gt dipole-dipole gt London Dispersion
• Bonds interactions forces

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Water and Hydrogen Bonds

• Hydrogen bonds are extremely important in
  determining the properties of water and
  biological molecules such as proteins.
• The water molecule has a bent shape (105) and is
  considered to be polar and the universal solvent.
• The attraction in water results from the
  intermolecular hydrogen bonds.

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Water and Hydrogen Bonds

• Surface tension the inward force, or pull that
  tends to minimize the surface area of a liquid
• - this surface tension tends to hold a drop
  of liquid in a spherical shape
• The higher the surface tension, the more nearly
  spherical is the drop of that particular.

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Water and Hydrogen Bonds

• Because of hydrogen bonding, water absorbs a
  large amount of heat as it evaporates or
  vaporizes.
• The hydrogen bonds must be broken before water
  changes from the liquid to vapor state.
• Vapor Pressure the force exerted due to the gas
  above the liquid

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Water and Hydrogen Bonds

• Boiling Point occurs when the temperature at
  which the vapor pressure of the liquid is just
  equal to the external pressure.
• Boiling leads to evaporation of a liquid. In the
  case of water, hydrogen bonds break in order for
  the liquid to vaporize.