Ch. 12---Chemical Bonding

1
Ch. 12---Chemical Bonding

• Covalent Bonds
• ____________ electrons between two atoms in
  order to fill the outer energy level (or shell)
• Each bond involves the sharing of _____
  _________ of electrons.
• Single Bonds __ e-s Double Bonds __ e-s
  Triple Bonds__ e-s

Sharing
one pair
2
4
6
2

• Ways to Represent Covalent Bonds in Compounds
• Circled Arrows for bonds.
• a) H2 H ___
• 1s
• H ___
• 1s
• b) F2 F ___ ___ ___ ___
• 2s 2p
• F ___ ___ ___ ___
• 2s 2p

?
?
?
?
?
?
?
?
?
?
?
?
?
?
?
?
3

• c) NH3 N ___ ___ ___
  ___
• 2s
  2p
• d) H2O O ___ ___ ___ ___
  
• 2s
  2p

?
?
?
?
?
?
?
?
H ___ 1s
H ___ 1s
H ___ 1s
?
?
?
?
?
?
?
H ___ 1s
?
H ___ 1s
4

• Ways to Represent Covalent Bonds in Compounds
• (2) Dots for bonds. (Lewis Structures)
• a) H2 H H ?
• b) F2 F F ?
• c) N2 N N ?
  (triple bond)
• d) NH3

HH
..
..
..
..
FF
..
..
..
..




NN


..
HNH
..
H
5

• Ways to Represent Covalent Bonds in Compounds
• (3) Lines for bonds.
• a) H2
• b) F2
• c) N2
• d) NH3
• e) H2O

HH
..
..
FF
..
..
NN

HNH
?
H
..
HO
?
H
6

• Coordinate Covalent Bonds
• Both of the electrons that make the bond come
  from the ________ _______________ .
• Example CO (carbon monoxide)

same
element
?
C ___ ___ ___ ___
2s 2p
?
?
?
?
?
O ___ ___ ___ ___
2s 2p
?
?
?
?
..
C O
..
..
Two of the bonds are normal, and the third bond
is a coordinate covalent bond.

C O

?
7

• Carbons Hybrid Orbital
• C ___ ___ ___ ___ (Before)
• 2s 2p

?
?
?
?
C ___ ___ ___ ___ (After 4
covalent bonds available)
?
?
?
?
?
2sp3
Practice Problem Draw CH4 using arrows, dots,
and lines for bonds.
?
?
?
?
C ___ ___ ___ ___
2sp3
H ___ 1s
H ___ 1s
?
?
?
?
H ___ 1s
H ___ 1s
8
The 7 Diatomic Elements

• Some elements will covalently bond to themselves
  to form a molecule composed of ____ atoms.
• These elements are never found in nature as
  single atoms. Instead, they will be bonded as a
  ________ when they are in the _________________
  state.
• The 7 diatomic elements are the gases H, O, N,
  and all of the _________________, (Group 7A).
• H2, O2, N2, Cl2, Br2, I2, F2
• HONClBrIF

two
pair
elemental
halogens
9
(No Transcript)
10
Air contains N2 and O2 molecules.
11
The decomposition of two water molecules
12

• Octet Rule
• Atoms want ___ e-s in their outer shell when
  forming compounds.
• This will mean ___ dots around them all
  together. This is the stable e- configuration of
  a __________ _______!
• Important exception Hydrogen only needs __ to
  be full (like He).
• Other Exceptions
• PCl5 (___ e-) SF6 (___ e-)
  BF3(___e-)

8
8
noble gas
2
10
12
6
13
Resonance

• Resonance is the ability to draw 2 or more
  different e- dot notations that obey the octet
  rule.
• Examples O3 (ozone) and SO2

Practice Problem Draw the resonance structures
for CO3-2.
14
VSEPR Theory Molecular Shapes

• Most shapes are based on a __________________.
• Examples CH4 CCl4
• Removing the top of the tetrahedral makes the
  ________________ shape.
• Examples NH3 PCl3

tetrahedral
pyramidal
15
VSEPR Theory Molecular Shapes

• Removing one side of the pyramid makes the
  _____________ shape.
• Examples H2O H2S
• If there are only two atoms bonded, it is
  ______________.
• Examples O2 HCl CO2 (linear because of its
  double bonds.)

bent
linear
16
VSEPR Theory Molecular Shapes

• Another we will need to know is called trigonal
  planar. Trigonal means that the central atom
  is bonded to ___ other atoms. Planar means
  that the 3 atoms all lie in the same
  ______________.
• Example BF3
• (Notice that Boron will only have ___ e-s
  around it. The missing pair of electrons will
  make it planar instead of ________________.)

3
plane
6
pyramidal
17
VSEPR Theory Molecular Shapes

• Finally, the last 2 shapes occur when there are
  5 or 6 regions of electrons are around the
  central atom. (These molecules are also
  exceptions to the octet rule!)
• ___________________ (5 electron domains)
• __________________ (6 electron domains)
• Examples PCl5 and SF6

Trigonal bipyramid
Octahedral
18
Polar and Nonpolar Bonds

• Even though the electrons in a covalent bond are
  shared, sometimes the attraction for the bonded
  pair, (the _____________________), is uneven.
  This gives rise to 3 bond types.
• nonpolar covalent bonds ____________ sharing
  of the e- pair
• polar covalent bonds ________________ sharing
  of the e- pair
• ionic bonds a ___________ of e-s from the
  metal to the nonmetal
• How To Determine the Bond Type
• Bond type is based on the electronegativity
  _____________ between the two bonded atoms.
• (See p.403 for electronegativity values.)

electronegativity
equal
unequal
transfer
difference
19
(No Transcript)
20
Figure 12.4 The three possible types of bonds.
nonpolar
polar
ionic
21

• How To Determine the Bond Type
• 0 to 0.4 ______________ covalent bond
• 0.5 to 2.0 _____________ covalent bond
• Above 2.0 _______________ bond
• Practice Problems Determine the type of bond
  that forms between the atoms in the following
  compounds.
• a) CO2 b) NaCl c) CH4

nonpolar
polar
ionic
2.5 3.5
0.9 3.0
2.5 2.1
1.0 polar covalent
2.1 ionic
0.4 nonpolar covalent
22
Polarity of Molecules

• One side is slightly (__) and the other side is
  slightly (__).
• Polar molecules are also known as
  _______________.
• Polarity depends on the __________ and symmetry
  of the molecule.
• symmetrical molecules (looks the same on all
  sides) ___________
• asymmetrical molecules ___________
• Polar molecules are moved by ____________
  charges. (DEMO!)

dipole
shape
nonpolar
polar
static
23
Molecular Polarity

• Practice Problems Determine if the following
  molecules are polar or nonpolar based on their
  shape.
• a) CH4 b) NH3 c) H2O d) HCl
  e) BF3

symmetrical tetrahedral
asymmetrical pyramid
asymmetrical bent
asymmetrical linear
symmetrical trigonal planar
nonpolar
nonpolar
polar
polar
polar
Dipole of NH3
Dipole of H2O
24
Bond Dissociation Energy

• This is the energy needed to ___________ the
  bond.
• Generally, the longer the bond, the _____ energy
  it takes to break it.
• Single bonds take ________ energy to break than
  double bonds and triple bonds require the
  _________ energy to break.
• When bonds form, energy is _____________.
• (Breaking bonds requires the addition of energy.)

break
less
less
most
released
25
Bond Dissociation Energy
26
Intermolecular Attractions

• The __________ attractions between one molecule
  and another are called _______ ______ ________
  forces.
• They cause gas particles to stick together and
  _______________ at low temperatures.

weak
Van der Waals
condense
27
Dispersion Forces

• There are two types of intermolecular forces
• (1) ____________________ forces (the weaker
  type)
• caused by random _______________ motion
• generally _____________ with ________ electrons
  in the molecule

Dispersion
electron
stronger
more
- exist between all types of
molecules - This force causes Br2 to
be a liquid and I2 to be a solid
at room temperature.
28
Dipole Interaction Forces

• (2) ____________ interactions (the stronger
  force)
• caused by the attraction of the (__) side of one
  polar molecule and the (__) side of a different
  polar molecule

Dipole


29
Hydrogen Bonds

• ________________ Bonds are a special type of
  dipole interaction.
• They occur between the hydrogen of one polar
  molecule and the ____, ___ or ___ of another
  polar molecule.

Hydrogen
N O F
Hydrogen Bonding in Water
30
Hydrogen Bonds

• The ladder rungs in a DNA molecule are hydrogen
  bonds between the base pairs, (AT and GC).

31
Hydrogen Bonds in DNA
32
Ionic Bonding Ionic Compounds

• Ionic Bonds
• Form when ___________ transfer their
  _____________ electrons to a _______________.
• The forces of attraction between the
  ____________ () and the _____________ (-) bind
  the compound together.
• How to Represent an Ionic Bond
• Electron Configuration
• Na 1s2 2s2 2p6 3s1
• Cl 1s2 2s2 2p6 3s2 3p5

metals
valence
nonmetal
cation
anion
Na ___
3s
?
Cl ___ ___ ___ ___
3s 3p
?
?
?
?
?
?
?
33
How to Represent an Ionic Bond

• 2) Electron Dot Notations
• Na Cl ?
• Practice Problems (1) Draw the electron dot
  notation for the formation of an ionic compound
  between sodium and oxygen.
• (2) Draw the electron configuration notation for
  the formation of an ionic compound between
  magnesium and fluorine.

Na1
Cl -1
34
(No Transcript)
35

• Practice Problems
• 3) a) Draw the electron dot notation for a
  potassium atom.
• b) Draw the electron dot notation for a
  potassium ion.
• (4) a) Draw the electron dot notation for a
  sulfur atom.
• b) Draw the electron dot notation for a sulfur
  ion.

K
K1
S
S -2
36
Properties of Ionic Compounds and Covalent
Molecules

• Ionic
• ?_______________ of electricity when dissolved
  water or melted.
• ?formed between __________ and _________________
• ?have _________ melting points
• ?usually ________ soluble in water
• form ___________________ solids

Conductors
metals
nonmetals
high
(dissolved salt)
very
ionic crystalline
37
Figure 15.1-- Polar water molecules interacting
with positive and negative ions of a salt.
38
Crystalline Patterns
39
Properties of Ionic Compounds and Covalent
Molecules

• Molecular
• ________________ of electricity
• formed between two _______________
• usually have ________ melting points
• solubility in water _______ (polar dissolve
  nonpolar insoluble)
• For a compound to to conduct electricity it must
  have
• (1) Charged Particles (________)
• (2) Particles Free to Move (___________ or
  __________ phase)

Insulators
nonmetals
low
varies
ions
liquid
aqueous
40
Demonstration
great conductor
good conductor
PureH2O
poor conductor
nonconductor