Bonding

1
Bonding
Chemical Bond

• A mutual electrical attraction between the nuclei
  and valence electrons of different atoms that
  bond the atoms together.

2
Types of Bonding

• Ionic Bond
• - results from a transfer of electrons. Held
  together by electrostatic force. (Opposites
  attract.)
• Covalent bond
• results from the sharing of electrons between
  two atoms.
• Non-polar covalent
• bond in which electrons are equally shared.
• Polar covalent
• bonds that have an uneven distribution of
  charge due to unequal attraction of the shared
  electrons.

3

• Bonds between two unlike atoms are never
  completely ionic and are rarely completely
  covalent.
• The degree to which bonds are ionic or covalent
  depends on the electronegativity differences of
  the bonded atoms.
• Bonding can be thought of as a tug-of war for
  electrons.
• Electronegativity is an elements strength.
• (Table p. 151).

4
Range of Bond Characteristic
Based on Electronegativity Difference
Covalent
Ionic
0.3
1.7
Nonpolar Covalent
Polar Covalent
5
Practice with Bond Type
1. What type of bond forms between Hydrogen and
Chlorine? HCl Since it is electronegativity
difference, we will _____________.
subtract
Electronegativities H 2.1 Cl 3.0
Polar Covalent!
6
Practice with Bond Type
2. What type of bond forms between Hydrogen and
Fluorine? HF
Electronegativities H 2.1 F 4.0
Ionic!
7
Practice with Bond Type
3. What type of bond forms between Hydrogen and
Hydrogen? H2
Electronegativity H 2.1
Nonpolar Covalent!
8
Covalent Bonding
9
Important Terms of Covalent Bonding

• Molecule - resulting particle when two or more
  atoms bond covalently.
• Diatomic molecule - molecule consisting of two
  atoms.
• Seven elements occur as diatomic molecules in
  the natural state.
• Bromine, Iodine, Nitrogen, Chlorine, Hydrogen,
  Oxygen, Fluorine
• Br2 I2 N2 Cl2 H2 O2 F2
• Single bond - covalent bond produced by the
  sharing of one pair of electrons.
• Double bond - sharing of two pairs of electrons.
• Triple bond - sharing of three pairs of
  electrons.
• Bond length - the average distance between the
  nuclei of two bonded atoms.
• Bond energy - energy required to break a chemical
  bond and form neutral atoms.

10
Lewis Structures

• Structural formula - indicates the kind, number,
  arrangement, and bonds of the atoms in a
  molecule.
• Atomic symbols represent inner-shell electrons
  and nuclei
• Dashes between two atomic symbols represent
  shared electron pairs in covalent bonds
• Dots adjacent to only one atomic symbol
  represent unshared or lone electrons.
• - represents a bond (2 electrons)
• ? represents an unshared electron
• Central atom is the least electronegative atom
  (furthest to the left on the periodic table)
  (Except hydrogen)
• All atoms need to have 8 electrons in their
  outer level to be stable except H
• H needs 2electrons in its outer shell.

11
Drawing Lewis Structures

• Step 1 List all the atoms in the compound.
• Example CH4

C
H
H
H
H
12
Drawing Lewis Structures Cont.

• Step 2Find the total of electrons needed to
  have a complete outer shell.
• (Remember All elements need 8 except H, which
  needs only 2).

Need
C
8
H
2
H
2
H
2
H
2
16
13
Drawing Lewis Structures Cont.

• Step 3 Find the total number of valence
  electrons that the elements have available.
• (Use PT).

Need
Have
C
4
8
H
1
2
1
H
2
H
1
2
H
1
2
16
8
14
Drawing Lewis Structures Cont.

• Step 4 Subtract the HAVES from the NEED.
  This is the number of electrons that must be
  SHARED.

Need
Have
C
4
8
H
1
2
1
H
2
H
1
2
H
1
2
Shared e-
16
8
8
15
Drawing Lewis Structures Cont.

• Step 5 Find the number of bonds.
• Since two electrons are shared per bond, divide
  the number of SHARED electrons by 2 to get the
  number of bonds.

Need
Have
C
4
8
H
1
2
1
H
2
H
1
2
H
1
2
Shared e-
16
8
8
4 bonds
2
16
Drawing Lewis Structures Cont.

• Step 5 Draw the structure.
• Put the atom furthest left on the periodic table
  (Except H) in the center. Fill in the bonds
  (dashes) and lone e- (dots) to ensure that each
  atom has the number of electrons it needs.

H
C
H
H
H
COUNT Electrons for all atoms and the entire
molecule!
17
Practice Draw Lewis structures for the
following molecules.

• 1) CH3I
• 2) Silicon dioxide (SiO2)
• 3) Ammonia (NH3)
• 4) Iodine monochloride (ICl)
• 5) Water

18
Lets do the first few together!
1) CH3I
Need
Have
C
4
8
H
1
2
1
H
2
H
1
2
I
7
8
Shared e-
22
14
8
4 bonds
2
19
CH3I Cont.
Structure
COUNT Electrons for all atoms and the entire
molecule!
H
C
I
H
H
H All have two electrons (1 bond) C- Has 8
electrons (4 bonds) I - Has 2 electrons (1
bond). I needs 6 unshared electrons Molecule has
14 total electrons (8 shared 6 unshared!)
20
2) Silicon dioxide (Si02)
Need
Have
Si
4
8
Si
O
O
O
6
8
6
O
8
Shared e-
24
16
8
4 bonds
2
21
3) Ammonia (NH3)
Need
Have
H
N
5
8
H
1
2
N
H
1
H
2
H
1
2
Shared e-
H
14
8
6
3 bonds
2
22
4) Iodine monochloride (ICl)
Need
Have
I
7
8
Shared e-
Cl
7
8
Cl
I
16
14
2
1 bond
2
23
5) Water
Need
Have
H
1
2
O
1
H
2
H
H
O
6
8
Shared e-
12
8
4
2 bonds
2
24
VSEPR THEORY

• Valence Shell, Electron Pair Repulsion theory
• Determines the shape of the molecule.

unshared unshared repulsion
unshared shared repulsion
shared shared repulsion
gt
gt
O
H
H
25
Guide for Determining Shapes
1
Linear
NA
No
Linear
2
Yes
Bent

No
Trigonal Planar
3
Yes
Trigonal Pyramidal
4
No
Tetrahedral
26
Molecular Shapes
Look at CENTRAL atom when determining shape!
Linear

• 1. _________________
• 2. _________________
• 3. _________________
• 4. _________________
• 5. _________________

Bent
Trigonal planar
Trigonal pyramidal
Tetrahedral
27
IONIC COMPOUNDS

• Composed of positive and negative ions combined
  so that the positive and negative charges are
  equal.
• Formula Unit
• the simplest collection of atoms from which a
  compounds formula can be established. Ex. 1
  formula unit of sodium chloride (NaCl) is one
  sodium ion plus one chloride ion. The formula
  unit of a compound depends on the charges of the
  ions combined.
• Ex. magnesium chloride MgCl2

28
Properties of Ionic Bonds

• High melting points
• Hard, brittle crystalline solids.
• Good conductors of electricity in molten form or
  when dissolved in water.

29
Metallic Bonding

• Sharing of delocalized electrons. Results from
  the attraction of positive ions and surrounding
  mobile electrons.