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CHEMICAL BONDING

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Title: CHEMICAL BONDING


1
CHEMICAL BONDING
According to the types of bonds contained in a
molecule, the physical properties including
melting point, hardness, electrical and thermal
conductivity and solubility are determined.
  • Cocaine

2
Chemical Bonding
  • Problems and questions
  • How is a molecule or polyatomic ion held
    together?
  • Why are atoms distributed at strange angles?
  • Why are molecules not flat?
  • Can we predict the structure?
  • How is structure related to chemical and physical
    properties?

3

What is a chemical bond?
  • Attraction between two or more atoms
  • Interaction between valence electrons

4

Why do atoms attract?
  • To become more stable!
  • What is the magic
  • number?

5

Octet Rule?
  • Atoms want to have 8 valence electrons to be
    more stable
  • Atoms will combine when the compound formed is
    more stable than the separate atoms
  • Which atoms need less than 8 valence e- to be
    stable?

6
How is Stability Reached?
  • Atoms will either gain, lose, or share electrons
    to achieve stability (noble gas configuration)

7
The type of bond can usually be calculated by
finding the difference in electronegativity of
the two atoms that are going together.
8
Common Ions
  • Metals form positive ions and
  • Nonmetals form negative ions

1
-3
2
-1
3
-2
/-4
Varies Can be 1, 2, or 3
9
Electron Distribution in Molecules
  • Electron distribution is depicted with Lewis
    (electron dot) structures
  • This is how you decide how many atoms will bond
    covalently! (In ionic bonds, it was decided
    with charges)

10
Lewis Dot Structures
  • Lewis dot structures are a shorthand to represent
    the valence electrons of an atom.
  • The structures are written as the element symbol
    surrounded by dots that represent the valence
    electrons.
  • The Lewis structures for the elements in the
    first two periods of the periodic table are shown
    below.

11
Lewis Dot Structures
Example Write the electron dot symbol for
phosphorus. Answer P Phosphorus Group 5A,
therefore, there are 5 valence electrons.
12
Lewis Dot Structures
Example Problem Use Electron Dot Symbols to
represent the formation of Magnesium Fluoride
from atoms of Mg and F. Mg Group 2A F
Group 7A Mg F2 ? MgF2
13
Molecule H2O  NH3 CH4 
Number of each kind of atom in molecule O 1 H 2 N 1 H 3 C 1 H 4
Valence electrons for each atom O 6 H 1 N 5 H 1 C 4 H 1
Total number of valence electrons O 1 x 6 6 H 2 x 1 2 6 2 8  N 1 x 5 5 H 3 x 1 3 5 3 8  C 1 x 4 4 H 4 x 1 4 4 4 8
Arrangement of dots
14
Electronegativity
Francium has the LOWEST electronegativity. Fluorin
e has the HIGHEST electronegativity
15
Review of Chemical Bonds
  • There are 3 forms of bonding
  • _________complete transfer of 1 or more
    electrons from one atom to another (one loses,
    the other gains) forming oppositely charged ions
    that attract one another
  • _________valence electrons shared between atoms
  • _________ holds atoms of a metal together

Most bonds are somewhere in between ionic and
covalent.
16
Electronegativity Difference
  • If the difference in electronegativities is
  • gt2.0 Ionic (NaCl 3.0 0.8 2.2)
  • 1.0-2.0 Very Polar Covalent (HF 1.9 H2O
    1.4)
  • 0.4-1.0 Moderately Polar Covalent (HCl
    0.9)
  • 0.0 to 0.4 Non-Polar Covalent (H2 0.0)

Example NaCl Na 0.8, Cl 3.0 Difference is
2.2, so this is an ionic bond!
17
SUMMARY OF BOND TYPES
18
Why is it called ionic?
  • Sodium loses one electron to form a positive ion
    while chlorine gains one electron to form a
    negative ion.
  • Electron Transfer

19
Ionic Bonding
  • A bond between a metal and a nonmetal


20
Ionic Bonds
  • All those ionic compounds were made from ionic
    bonds. Weve been through this in great detail
    already. Positive cations and the negative
    anions are attracted to one another (remember the
    Paula Abdul Principle of Chemistry Opposites
    Attract!)

Therefore, ionic compounds are usually between
metals and nonmetals (opposite ends of the
periodic table).
21
Properties of Ionic Compounds
  • Compounds containing a metal and a nonmetal
  • Have high melting points
  • Most dissolve in water
  • When dissolved in water, they conduct electricity

22
SUMMARY OF BOND TYPES
  • NONPOLAR COVALENT BOND Equal sharing of
    electrons, because the nuclear attraction for the
    electron pair is EQUAL.
  • POLAR COVALENT BOND Unequal sharing of
    electrons, because one atom has a greater nuclear
    attraction for the electron pair than the other
    atom.
  • IONIC BOND The metal donates its electrons to
    obtain a positive charge and the nonmetal accepts
    the electrons from the metal to obtain a negative
    charge.

23
Covalent Bonding
  • Bond formed between 2 nonmetals

24

Why is it called covalent?
  • Covalent essentially means sharing electrons

25

Properties of Covalent Compounds
  • Compounds of 2 nonmetals
  • Have low melting points
  • Most do not dissolve in water
  • Do not conduct electricity in solution

26

Metallic Bonding
  • Bond formed between metal atoms
  • Valence electrons move freely

27
Properties of Metallic Bonds
  • Metals have high melting points
  • Metallic bonds are strong
  • Good conductors of heat and electricity

28
Bond and Lone Pairs
  • Valence electrons are distributed as shared or
    BOND PAIRS and unshared or LONE PAIRS.



This is called a LEWIS structure.
29
Bond Formation
  • A bond can result from an overlap of atomic
    orbitals on neighboring atoms.




Overlap of H (1s) and Cl (2p)
Note that each atom has a single, unpaired
electron.
30
Review of Valence Electrons
  • Remember from the electron chapter that valence
    electrons are the electrons in the OUTERMOST
    energy level thats why we did all those
    electron configurations!
  • B is 1s2 2s2 2p1 so the outer energy level is 2,
    and there are 21 3 electrons in level 2.
    These are the valence electrons!
  • Br is Ar 4s2 3d10 4p5How many valence
    electrons are present?

31
Review of Valence Electrons
  • Number of valence electrons of a main (A) group
    atom Group number

32
Steps for Building a Dot Structure
  • Ammonia, NH3
  • 1. Decide on the central atom never H. Why?
  • If there is a choice, the central atom is atom
    of lowest affinity for electrons. (Most of the
    time, this is the least electronegative atomin
    advanced chemistry we use a thing called formal
    charge to determine the central atom. But thats
    another story!) Therefore, N is central on this
    one
  • 2. Add up the number of valence electrons that
    can be used.
  • H 1 and N 5
  • Total (3 x 1) 5
  • 8 electrons / 4 pairs

33
Building a Dot Structure
  • 3. Form a single bond between the central atom
    and each surrounding atom (each bond takes 2
    electrons!)

4. Remaining electrons form LONE PAIRS to
complete the octet as needed (or duet in the case
of H).
3 BOND PAIRS and 1 LONE PAIR.
Note that N has a share in 4 pairs (8 electrons),
while H shares 1 pair.
34
Building a Dot Structure
  1. Check to make sure there are 8 electrons around
    each atom except H. H should only have 2
    electrons. This includes SHARED pairs.

6. Also, check the number of electrons in your
drawing with the number of electrons from step 2.
If you have more electrons in the drawing than
in step 2, you must make double or triple bonds.
If you have less electrons in the drawing than in
step 2, you made a mistake!
35
Carbon Dioxide, CO2
  • 1. Central atom
  • 2. Valence electrons
  • 3. Form bonds.

C 4 e-O 6 e- X 2 Os 12 e-Total 16 valence
electrons
This leaves 12 electrons (6 pair).
4. Place lone pairs on outer atoms.
5. Check to see that all atoms have 8 electrons
around it except for H, which can have 2.
36
Carbon Dioxide, CO2
C 4 e-O 6 e- X 2 Os 12 e-Total 16 valence
electrons How many are in the drawing?
6. There are too many electrons in our drawing.
We must form DOUBLE BONDS between C and O.
Instead of sharing only 1 pair, a double bond
shares 2 pairs. So one pair is taken away from
each atom and replaced with another bond.
37
Double and even triple bonds are commonly
observed for C, N, P, O, and S
H2CO
SO3
C2F4
38
Now You Try One!Draw Sulfur Dioxide, SO2
39
Violations of the Octet Rule(Honors only)
  • Usually occurs with B and elements of higher
    periods. Common exceptions are Be, B, P, S, and
    Xe.

Be 4 B 6 P 8 OR 10 S 8, 10, OR 12 Xe 8, 10,
OR 12
40
MOLECULAR GEOMETRY
41
MOLECULAR GEOMETRY
Molecule adopts the shape that minimizes the
electron pair repulsions.
  • VSEPR
  • Valence Shell Electron Pair Repulsion theory.
  • Most important factor in determining geometry is
    relative repulsion between electron pairs.

42
Some Common Geometries
Linear
Tetrahedral
Trigonal Planar
43
VSEPR charts
  • Use the Lewis structure to determine the geometry
    of the molecule
  • Electron arrangement establishes the bond angles
  • Molecule takes the shape of that portion of the
    electron arrangement
  • Charts look at the CENTRAL atom for all data!
  • Think REGIONS OF ELECTRON DENSITY rather than
    bonds (for instance, a double bond would only be
    1 region)

44
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45
Other VSEPR charts
46
Structure Determination by VSEPR
  • Water, H2O

The electron pair geometry is TETRAHEDRAL
2 bond pairs 2 lone pairs
The molecular geometry is BENT.
47
Structure Determination by VSEPR
  • Ammonia, NH3
  • The electron pair geometry is tetrahedral.

The MOLECULAR GEOMETRY the positions of the
atoms is TRIGONAL PYRAMID.
48
Bond Polarity
  • HCl is POLAR because it has a positive end and a
    negative end. (difference in electronegativity)

Cl has a greater share in bonding electrons than
does H.
Cl has slight negative charge (-d) and H has
slight positive charge ( d)
49
Bond Polarity
  • This is why oil and water will not mix! Oil is
    nonpolar, and water is polar.
  • The two will repel each other, and so you can not
    dissolve one in the other

50
Bond Polarity
  • Like Dissolves Like
  • Polar dissolves Polar
  • Nonpolar dissolves Nonpolar
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