1. 5.1 Ionic Bonds: Chemical Bonding
The majority of elements occur in nature in chemical
combination with other elements in the form of compounds
Exceptions:
Noble Gases: He, Ne, Ar, Kr, & Xe
Some Metals: Cu, Ag, Au, Pt
O2, N2 & S8 occur in nature as diatomic and
polyatomic molecules
The electrons of the interacting atoms are involved in compound
formation in one of two ways:
Transferring Electrons between Atoms - Ionic Compounds
Sharing Electrons between Atoms - Covalent Compounds
It is the Electronic Structure of Each Element
that Gives Rise to the Chemistry (Reactivity) of
Each Element in the Periodic Table
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2. 5.1 Ionic Bonds: Ionic Bond Formation
The noble gases have stability (low reactivity) that is related
to the number and arrangement of electrons
Metallic Elements lose electrons to form ions with the
same number of electrons as the nearest noble gas
Non-Metallic elements gain electrons to form ions with
the same number of electrons as the nearest noble gas
The Formation of Sodium Chloride:
2Nao Cl NaCl
Sodium (11 e - ) loses 1 electron to form Na + (10 e - ) and
attains the stability of Neon (10 e - ).
Chlorine (17 e - ) gains 1 electron to form Cl - (18 e - ) and
attains the stability of Argon (18 e - ).
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3. 5.1 Ionic Bonds: - The Periodic Table
Non-Metals Gain Electrons
Metals Lose Electrons
Noble
Gases 5
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4. 5.2 Ionic Compounds: Ionic Compounds
Ionic Compounds: “ Ionic compounds are composed of ions.
An ion is any species (element or
compound) that bear a charge “
Binary Ionic Compound: “ One composed of just two elements;
formed when a metallic element
reacts with a non-metallic element “
Cation Formation: “ The metal loses a certain number of
electrons and becomes a positively
charged cation “
Cation: “ A positively charged ion “
Anion Formation: “ The non-metal gains a certain number of
electrons and becomes an anion “
Anion: “ A negatively charged ion “
In Effect, the Metal Atom Transfers Electrons
to the Non-Metal Atom
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5. 5.2 Ionic Compounds: Sodium Chloride (NaCl) Lattice Structure Form
2Nao (s) Cl2 (g) NaCl (s)
The Chemical Equation:
The Particulate (Atomic) View
Chlorine Ion ( Cl - )
Chlorine
Atom (Cl)
Sodium Ion ( Na + )
Na + and Cl - compound
Shown in the Crystal
Lattice Structure Form
Sodium Chloride
Crystal
Sodium Atom (Na)
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6. 5.3 Covalent Bonds
Covalent compounds form when elements share electrons,
which usually occurs between non-metallic elements.
Covalent Bond: “ In a covalent bond the pair of electrons is
mutually attracted by two nuclei.
The electrons are “ shared ” between the
two atomic elements “
H H H : H or H2
The electron from each
hydrogen atom is shared
between both hydrogen atoms
to give the helium noble
gas configuration ( He: )
Hydrogen
Atoms
Diatomic Hydrogen
Molecule
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7. Formation of a Covalent Bond between Two Hydrogen Atoms
H H H : H
It is the attraction of the bonding
electrons to two nuclei that holds
the nuclei together in a covalent bond
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8. 5.3 Covalent Bonds (cont)
Lewis Structures: * “ The element displayed with an electron dot
symbol to show the valence electrons “
* Also called Lewis symbols or Lewis diagrams.
Group
E.C.
No V. e-
Lewis
Symbol 1A
ns1 1
Li . 2A
ns2 2
. Be . 3A
ns2np1 3 .
. B . 4A
ns2np2 4 .
. C . 5A
ns2np3 5 .
. N : 6A
ns2np4 6 .
: O : 7A
ns2np5 7 .
: F :
. . 8A
ns2np6 8
. .
: Ne :
E.C = Electron Configuration
No V. e- = Number of Valence Electrons
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9. 5.3 Covalent Bonds (cont) . . The Reaction of Hydrogen and Fluorine:
H2 (g) F2 (g) HF (g)
Covalent Bond Formation from a Lewis Structure Viewpoint:
The octet
of electrons
around F is
complete
Vacant Orbital .
: F :
. .
H . +
H : F :
A single line
Is also used
to describe
the covalent
bond H
. .
F :
Each single bond (line)
is made up of 2 shared electrons
Octet Rule: “When atoms bond, they lose, gain, or share electrons
to attain a filled outer shell of eight (or two) electrons “
This completes the noble gas configuration
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10. 5.3 Covalent Bonds (cont) . Molecules with Covalent Bonds:
H Br H : Br
Hydrogen Bromide
. .
2 H H : O : H .
: O :
Water
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11. 5.3 Covalent Bonds (cont) . . Diatomic Molecules: . : F : + : F : F :
The lone pairs
are often not
shown
F . +
F F
Other Diatomic Molecules:
Cl . +
Cl Cl
Br . +
Br Br
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12. 5.3 Covalent Bonds (cont) . N : . N : Multiple Bonds - Diatomic Oxygen
Double Bonds are formed when 4 electrons (2 pairs of electrons)
are shared between two atoms .
: O : +
: O : : O :
Each oxygen
atom has a complete
octet of electrons
There is a double
bond between these
two oxygen atoms
O O
Diatomic Nitrogen
Triple Bonds are formed when 6 electrons (3 pairs of electrons)
are shared between two atoms .
. N : .
. N :
: N : : : N :
Each nitrogen
atom has a complete
octet of electrons +
There is a triple
bond between these
two nitrogen atoms
N N
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13. 5.3 Covalent Bond: Bond Length & Bond Strength
Triple bonds are shorter than double bonds
Double bonds are shorter than single bonds
Bond Lengths and Bond Energies
Bond Type
C C
C C
C C
N N
Bond Length (nm)
0.154
0.134
0.120
0.109
Bond Strength (kcal/mol) 83
146
200
225
Bond Strength:
Triple bonds are stronger than double bonds
Double bonds are stronger than single bonds
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14. 5.4 Shapes of Molecules: Electron Pair (EP) Repulsion
Molecular shape plays an important role in determining the
macroscopic properties of a chemical substance
Valence Shell Electron-Pair Repulsion (VSEPR) Theory
The electron pairs that are drawn in a Lewis structures,
distribute themselves around a particular atom so that they
are as far apart from each other as possible
These are the locations of lowest potential energy
This arrangement is called electron-pair (or region) geometry
This lowest energy condition also holds for regions of
electron density other than electron pairs (ie, double bonds)
Electron Pair (or region) Geometries
No. Electron Regions Geometry Electron-Pair Angles
Linear o
Trigonal Planar o
Tetrahedral o
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15. 5.4 Shapes of Molecules: Electron Pair Geometry
Two Electron Pairs (or Regions) Around an Atom
Linear Geometry
Electron Regions are 180 o
apart from each other
Three Electron Pairs (or Regions) Around an Atom
Trigonal Planar Geomety
Electron Regions are 120 o
apart from each other
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16. 5.4 Shapes of Molecules: Electron Pair Geometry (cont)
Four Electron Pairs (or Regions) Around an Atom
Tetrahedral Geometry
Electron Regions are o
apart from each other
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17. 5.4 Shapes of Molecules: Electron Pair Geometry (cont)
Linear
2 atoms on opposite sides of central atom
180° bond angles
Trigonal Planar
3 atoms form a triangle around the central atom
Planar
120° bond angles
Tetrahedral
4 surrounding atoms form a tetrahedron around the central atom
109.5° bond angles
180°
120°
109.5°
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18. 5.4 Shapes of Molecules: Molecular Shape
Linear Electron Pair Geometry
2 Regions of electrons around the central atom, both bonding
Or two atom molecule as trivial case ( N2, O2, etc )
Trigonal Electron Pair Geometry
3 Regions of electrons around the central atom
All Bonding = trigonal planar ( CO3 2- )
2 Bonding + 1 Lone Pair = trigonal bent ( O3 )
Tetrahedral Electron Pair Geometry
4 Regions of electrons around the central atom
All Bonding = tetrahedral
3 Bonding + 1 Lone Pair = trigonal pyramid ( NH3 )
2 Bonding + 2 Lone Pair = tetrahedral bent or V-shaped ( H2O )
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19. 5.5 Polar and Nonpolar Bonding
Nonpolar Bond: “ A bond in which bonding electrons are shared
equally is a nonpolar bond “
A bond between identical atoms ( H2 , F2, etc) is always nonpolar
Polar Bond: “ A bond in which bonding electrons are shared unequally “
Electrons
Distributed
Evenly
Unevenly 11
Chlorine attracts the
bonding electrons more
strongly than hydrogen
Chlorine is More Electronegative than Hydrogen
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20. 5.5 Polar and Nonpolar Bonding (cont)
Electronegativity: “ A measure of how strong an atom attracts
bonded electrons to itself.
A measure of how stingy an atom is for bonded electrons
Covalent bonding between unlike atoms results in unequal sharing of the electrons
One end of the bond has greater electron density than the other
The end with the larger electron density gets a partial negative charge
The end that is electron deficient gets a partial positive charge
H F •  
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21. The Polar Nature of Water
Covalent Bond Formation - 1) The Distribution of Bonding Electrons
In a covalent bond formed from different atoms
the sharing of electrons is not equal
One atom will prefer to have a greater share of the electrons
(one atom is more electronegative than the other)
Since the oxygen atom is more electronegative than hydrogen
the bonding electrons will spend more time around oxygen
This unequal sharing creates partially charged “ poles “
at the end of each O  H bond
Partial Negative
Charge on Oxygen
The Water Molecule O
H H
Partial Positive
Charge on Hydrogen
Electrons have Shifted the Average Position Nearer to Oxygen
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22. The Polar Nature of Water (cont)
Covalent Bond Formation - 1) The Distribution of Bonding Electrons (cont)
The bond polarity is represented by a polar arrow pointing
to the negative pole from a plus sign on the arrow
The Two
O  H Bonds
in Water
are Polar
Water has
a bent shape
Covalent Bond Formation - 2) The Overall Shape of the Water Molecule
The combined effects of
bent shape and polar bonds
makes water a polar molecule
Water is a
Polar Molecule
The Overall Vector
Component of Polarity
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