Bonding Unit
Today we will:
-Define Ionic, and Covalent Bonding
-Discuss ionic and covalent properties
-Learn to draw Lewis Structures
-Be Chemistry Match Makers
Ionic Compounds
•Ionic compounds are composed of both
metals and nonmetals. The bond that is
formed is based on electrostatic forces
between negatively(anion) and
positively(cation) charged ions.
•Ionic bonding occurs by the transfer of one or
more electrons from one atom to another
Properties
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usually solid at room temperature
have high melting points
usually do not conduct electricity as a solid
usually dissolve in water
usually conduct electricity when in solution or
molten state
Ionic Bonds
•An ion is an atom or group of atoms that have a charge. Atoms
normally have a neutral charge because most often they have the
same number of electrons and protons. They become ions by the
loss or addition of one or more electrons. This process is called
ionization.
•An ion that has more electrons than protons is called an anion,
and an ion that has fewer electrons than protons is called a
cation.
Ionic Bonds
•The interaction of ionic bonds is when atoms gain or lose
electrons until the outer shell of electrons is full and stable
with 8 electrons. This is part of the octet rule.
•Recall octet rule:
When atoms combine to form molecules they generally each
lose, gain, or share valence electrons until they attain or share
eight and reach a noble gas electron configuration
Ionic Bonds
•The number of electrons the atom gains, loses, or shares is called
its Valence.
•Nonmetals usually have four or more electrons in their outer
shell. To make their outer shell full, it’s easier(it takes less
ionization energy) for them to gain three or four electrons than to
lose four or five electrons.
•When you look at the metals, they usually have three or less
electrons in their outer shell. Opposite from nonmetals, it is easier
for metals to lose three or less electrons than to gain four or
more.
•Therefore it makes sense that metals and nonmetals bond
together easily.
Lewis Dot Structure
•In 1902 Gilbert Newton Lewis invented the valance bond
theory.
•Lewis came up with an easy way to represent electrons in the
outer shells of ions. His invention is called “Lewis Dot
Symbols”.
•Lewis structures are used to visualize the valence electrons of
elements. In the Lewis model, an element symbol is inside the
valence electrons of the s and p subshells of the outer ring.
Lewis Dot Structures
Lewis Dot Structure
1)Draw skeletal structure of compound showing what atoms are
bonded to each other. Put least electronegative element in the
center.
2)Count total number of valence e-. Add 1 for each negative
charge. Subtract 1 for each positive charge.
3)Complete an octet for all atoms except hydrogen
4)If structure contains too many electrons, form double and triple
bonds on central atom as needed.
Lewis Dot Structure
Example:
Sulfur has 16 electrons. You can show this by using the method
you learned earlier in the year [Ne]3s23p4. You can see from
using this method that there are 6 electrons in the outer shell of
sulfur. Sulfur can be represented in Lewis dots when you follow
these rules:
1. The element symbol is placed in the center.
2. Dots, representing electrons, are placed on all four sides of the element.
3. Each side can hold 2 electrons.
4. You must put an electron on each side (if possible) before filling up a side.
Lewis Dot Structure
This is what sulfur looks like according to the Lewis
Dot Chart:
S
Lewis Dot Structure
•Now it’s your turn to try and draw some elements using the
Lewis Dot Structure.
1. Potassium
2. Germanium
3. Phosphorus
4. Neon
5. Aluminum
Lewis Dot Structure
Covalent Compounds
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Covalent compounds are made up of two
nonmetals. A single covalent bond is formed
when a pair of electrons is shared between two
atoms
There are two types of covalent bonds: nonpolar covalent and polar covalent.
Properties
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simple molecular substances have low melting
and boiling points
larger more complex compounds will have
higher melting and boiling points
usually do not conduct electricity as a solid or
when molten or in solution
usually do not dissolve in water
Non-Polar Covalent Bonding
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Accounts for the bond that keeps two atoms of the
same element together. (Cl2, H2)
Atoms share electrons from ½ filled orbitals to achieve
noble gas configurations
Shared electrons are attracted to both nuclei, which
keeps atoms together
Electrons involved in bonding are called shared
electron pairs, ones that are not are called lone electron
pairs
Polar Covalent Bonds
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Account for the bonding found in HF
One electron from each. atom is shared but not equally due to
unequal attraction for shared electrons
The bond is referred to as polar because 2 poles are formed
(+ and -)
Electronegativity values allows us to determine which atom has
a greater pull
The atom with the greater electronegativity becomes the
negative end of the polar bond.
The atom with the lower electronegativity becomes the positive
end of the polar bond
Example-HF
e- poor
H
F
H
d+
e- rich
F
d-
Electronegativity
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Electronegativity is the tendency of an atom to draw
or attract the electrons in a bond toward itself
Electronegativity is like a game of tug-of-war, atom's
ability to pull determines what kind of bond it forms
To form a covalent bond, two or more atoms with
similar electronegativities will share electrons
Values fall between a low of 0.7 for Fr and a high 4.0
for F
The greater the difference in electronegativity the
more polar the bond.
Double Covalent Bonds
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Compounds sometimes share two pairs of
electrons and form a double bond. This often
occurs when two atoms of the same element
bond, but also occur between different
elements.
This is called Double Covalent bonds
Examples: O2, CO2
Triple Covalent Bond
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Same idea as single and double
Two atoms of the same element or two
different elements share three pairs of electrons
and form a triple bond
Example: N2
Coordinate Covalent Bond
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A covalent bond in which one of the atoms donates both
electrons
Properties are same as previous covalent bonds
Distinction is useful when keeping track of valence electrons
and when assigning formal charges.
Example formation of ammonium ion
Methods to Classify Bond Type
Method 1
• Subtract the
electronegativity of the
least electronegastive
atom from the most
• Divide the difference by
the greater value
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Covalent Bonds are in
the range of less than
0.17, polar covalent
bonds are between 0.17
and 0.45, and ionic bonds
are greater than 0.45
Method 2
• Covalent bonds=no
difference in
electronegativity
• Polar covalent
bonds=difference less
than 1.7
• Ionic Bonds= difference
of 1.7 or more.
Chemistry Match Makers
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Draw element
Research it
Make dating profile
½ class draw profiles
Get into group
Decide if it’s a match and draw it/classify
it/give reasons why
Report results