Transcript Chapter 9: Covalent Bonds

UNIT 9: Covalent Bonding
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What is a Covalent Bond?
• Covalent Bond –formed when two nonmetals
share pairs of valence electrons in order to
obtain the electron configuration of a noble
gas
• Molecule - formed when two or more atoms
bond covalently. (A molecule is to a covalent
bond as a formula unit is to an ionic bond.)
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Diatomic Molecules
• HOFBrINCl
Share electrons when they bond together
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Polyatomic Ions
• covalently bonded group of atoms, with a
charge
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Properties of Covalent Molecules
• Can exist as gases, liquids, or solids depending
on molecular mass and polarity
• Usually have lower MP and BP than ionic
compounds of the same mass
• Do not usually dissociate (break apart into
ions) in water
• Do not conduct electricity
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How to draw Lewis dot structures for covalent
molecules:
1.
2.
3.
Write the formula for the compound.
Count the total number of valence electrons.
Predict the location of the atoms:
a) If there is only 1 atom of an element, it is the central atom.
b) If carbon is present, it is ALWAYS the central atom.
c) The least electronegative atom is generally the central atom.
d) Hydrogen is NEVER the central atom.
4. Place one electron PAIR between the central atom and each ligand (side
atom) to “hook” the atoms together.
5. Dot the remaining electrons in pairs around the compound to complete
the octet. Start with the ligands.
6. Check that each atom has an octet. (H only needs a pair, not an octet.)
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Lewis Structures for Molecules
• Draw the Lewis dot structure for these
molecules:
– Hydrogen + Bromine (HBr)
– Carbon + Chlorine (CCl4)
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Writing Lewis Dot Structures - Covalent
Bonds
Bonding e- Pairs
Lone Pairs
(nonbonding electrons)
Covalent bonds
Exceptions to the octet rule:
• Molecules that have an odd # of valence electrons; ex.
NO2 has 17 total valence electrons and can’t form an
exact # of pairs
• Molecules with fewer than 8 electrons present; ex. BH3
where B only has and only needs 6 electrons
• Molecules with an expanded octet; ex. PCl5 where P
forms 5 bonds and SF6 where S forms 6 bonds
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Number of bonds
• Single Bonds - when one pair of e- is shared
between atoms
• Double bond – when atoms share 2 pairs of
valence electrons; ex. O2
• Triple bond – when atoms share 3 pairs of
valence electrons; ex. N2
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Describing bonds
• Sigma bond - the first bond between 2 atoms
– A single bond is a sigma bond.
• Pi bond - the second bond between 2 atoms
– A double bond consists of a sigma bond and a pi
bond.
– A triple bond consists of a sigma bond and two pi
bonds.
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Carbon can form single, double and
triple bonds with itself.
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Types of Bonds:
• Nonpolar covalent (also called pure
covalent or covalent) equal sharing of
electrons between atoms; occurs between
the atoms in a diatomic molecule
(HOFBrINCl) and between C and H; ex. CH4
• Polar covalent – unequal sharing of
electrons between atoms; occurs between
two nonmetals or a nonmetal and a
metalloid; ex. H2O
• Ionic – complete transfer of electrons;
occurs between m/nm, m/PAI, PAI/nm or
PAI/PAI; ex. NaCl
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Bond type
•
•
Difference in
electronegativity
values
Distance between
atoms on the
periodic table
Non-Polar Covalent
NPC
Small
Polar Covalent
PC
Ionic
I
medium
big
THIS IS A CONTINUUM. IT DESCRIBES THE “IONIC CHARACTER” OF THE
BOND.
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Practice:
• What type of bond exists in each of the
following?
• 1. HCl
2. CaO
• 3. H2O
• 4. Br2
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Why are molecular shapes
important?
The shape of a
molecule plays a very
important role in
determining its
properties.
Properties such as smell, taste, and
proper targeting (of drugs) are all the
result of molecular shape.
Molecular Shape
Lewis structures do not show how atoms in a
molecule are arranged in 3-dimensional
space.
Can you tell the
molecular shape of
CCl4 from its Lewis
structures?
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Molecular Shape
Example: Water is not linear!
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Atoms in a molecule try to spread out from one
another as much as possible to reduce the
charge repulsion between their outer electrons.
H
methane, CH4
H C H
Is this the farthest that the
hydrogens can get away from each
other?
H
90°
109.5°
science.howstuffworks.com
science.howstuffworks.com
This shape causes less repulsion between the bonding
pairs of electrons as the hydrogen atoms are farthest away
from each other.
Molecular Shape
apchemcyhs.wikispaces.com
Molecules adopt a geometry (shape) that minimizes e – e
repulsions. This occurs when e- pairs are as far apart as
possible.
Sample problem – molecular geometry
What is the shape of the following
molecules?
commons.wikimedia.org
http://winter.group.shef.ac.uk/
tetrahedral
trigonal planar
en.wikipedia.org
www.chriscrews.com
Bent or angular
Trigonal
pyramid
VSEPR
• Valence Shell Electron Pair Repulsion
• A theory that states that electron pairs repel
both bonding and non-bonding electrons
resulting in a stable (lowest-energy) 3dimensional geometry.
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TO DETERMINE MOLECULAR SHAPE
• Use VSEPR (valence shell electron pair repulsion) rules:
1) Draw the Lewis dot structure for the molecule
2) Identify the central atom
3) Count total # of electron pairs around the central atom
(stearic number)
4) Count # of bonding pairs of electrons (regions of electron
density) around the central atom
5) Count # of lone pairs of electrons around the central
atom; lone pairs take up a lot of space
6) Look at summary chart, identify shape
**shapes with no lone pairs are symmetrical
**shapes with lone pairs are assymmetrical
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Practice:
• Determine the shape.
1. NF3
2. SiCl4
3. H2O
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Water is a POLAR
molecule
The more electronegative atom will have a slight
negative charge, the area around the least
electronegative atom will have a slight positive
charge.
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Symmetric molecules tend to be nonpolar
Asymmetric molecules with polar bonds are polar
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Naming Binary Molecules
1.
Write the name of the
first element.
2. Change the nonmetal’s
ending to –ide.
3. Use prefixes to indicate
the number of each type
of atom.
*Exception-the first element
will never have the prefix
mono
–
–
–
–
–
–
–
–
–
–
1-mono
2-di
3-tri
4-tetra
5-penta
6-hexa
7-hepta
8-octa
9-nona
10-deca
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Practice
• Write the name for the following molecules:
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Writing formulas for molecules
• The prefixes tell you the subscript for each
atom.
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Practice
• Write the formulas for the following
molecules:
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Empirical formula
• Empirical - To be derived from observation,
experiment, or data.
• Empirical formula - the simplest whole number
ratio between two (or more) elements
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Steps to determine the empirical
formula of a compound:
1. Determine the mass of each element in the sample.
2. Divide the mass of each element by the molar mass
(from the PT) to determine the number of moles of each
element. Round to the thousandths (._ _ _ )!
3. Divide the # of moles of each element by the smallest #
of moles. This is the mole ratio for each element in the
compound.
4. If your answers to step 3 are whole numbers, these are
written as the subscripts.
5. If your answers to step 3 are NOT whole numbers,
multiply by 2, 3, or 5 to obtain a whole number if
increments of 0.5, 0.3 or 0.2 are given, respectively.
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example
What is the empirical formula for a sulfur oxide
compound containing 50% sulfur and 50% oxygen?
Step 1: Since % means “parts per hundred”, assume
we are working with a 100 g sample. That means
we have 50 g of sulfur and 50 g of oxygen.
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Step 2: Use dimensional analysis to convert
grams to moles
50 g S __1 mol_ = 1.558 moles S
32.1 g
50 g O _1 mol_
16.0 g
Label
Properly!!
= 3.125 moles O
Round to thousandths (._ _ _)
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Step 3: Divide by the smallest number of
moles to obtain a mole ratio.
1.558 moles S
=1S
1.558 moles S
3.125 moles O
Label
Properly!!
=2O
1.558 moles S
So, we have 1 S for every 2 O.
These numbers become the subscripts and the formula is SO2
In our example, we did not need the 4th step since the ratio came out to a whole
number.
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Another example
A compound contains 54.1 g of Mg and
45.9 g of P. Determine the compound’s
empirical formula.
Note: This time, we already have the number of
grams so we can skip to step 2.
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Step 2: Use Dimensional Analysis to
convert grams to moles.
54.1 g Mg___1 mol_ = 2.226 moles Mg
24.3 g
Label
45.9 g P __1 mol_ = 1.481 moles P Properly!!
31.0 g
Round to thousandths place (._ _ _ )
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Step 3: Divide by the smallest number of
moles to obtain a mole ratio.
1.482 moles P
=1P
1.482 moles P
2.226 moles Mg
Label
Properly!!
= 1.5 Mg
1.482 moles P
Notice, the bottom answer did not come out to a
whole number this time.
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Skip to step 5 since answers are not whole numbers.
Step 5: Multiply answers from step 3 so that you get
whole numbers.
We had 1.5 Mg and 1 P
0.5 = ½
flip it and you have your scale factor, 2.
1.5 x 2 = 3 Mg and 1 x 2 = 2 P.
The ratio did not change, it is just a whole number ratio
now.
So, we have 3 Mg for every 2 P or the formula Mg3P2
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Now that you know the steps, here is a
jingle to make them easier to remember:
Percent to mass
step 1
Mass to mole
step 2
Divide by small
step 3
Multiply ‘til whole steps 4 and 5
Note: You may not need all of the steps. 
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Molecular formula
 A formula that is reducible.
 It is a multiple of an empirical formula.
Ex. Can C8H12 be reduced?
 Of course, it’s divisible by 4. So, dividing by
4 reduces the formula to C2H3.
 C8H12 is the molecular formula.
 C2H3 is the empirical formula.
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This template can help you organize your information and find what
you are missing.
Empirical formula
molecular formula
Mass of empirical formula
Mass of molecular formula
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Ex. The molar mass of a molecular formula is 283.88 g/mole and
it’s empirical formula is P2O5. Determine the molecular formula.
 Draw your chart and fill in the info from the
problem.
P2O5
?
141.943 g/mole
283.88 g/mol
Now, divide the molar mass of the MF by the molar mass of the EF.
(283.88 g/mole)/(141.943 g/mole) = 2. Scale factor is 2.
Multiply the subscripts in the EF by 2 and the MF is… P4O10
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practice
A compound is made from 2.00 g carbon, 0.335 g
hydrogen, and 2.66 g oxygen. Its molar mass is
90.0 g/mole. Determine the molecular formula.
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