UNIT 9: Covalent Bonding

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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

• Write the formula for the compound.
• Count the total number of valence electrons.
• Predict the location of the atoms
• If there is only 1 atom of an element, it is the
  central atom.
• If carbon is present, it is ALWAYS the central
  atom.
• The least electronegative atom is generally the
  central atom.
• Hydrogen is NEVER the central atom.
• Place one electron PAIR between the central atom
  and each ligand (side atom) to hook the atoms
  together.
• Dot the remaining electrons in pairs around the
  compound to complete the octet. Start with the
  ligands.
• 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

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Exceptions to the octet rule

• Molecules that have an odd of valence
  electrons ex. NO2 has 17 total valence electrons
  and cant 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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This is a continuum. It describes the ionic
character of the bond.
  Bond type   Non-Polar Covalent Polar Covalent Ionic NPC PC I
Difference in electronegativity values Distance between atoms on the periodic table   Small medium big
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Practice

• What type of bond exists in each of the
  following?
• 1. HCl2. CaO
• 3. H2O
• 4. Br2

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Why are molecular shapes important?
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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.
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Molecular Shape
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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?
www.mikeblaber.org
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Molecular Shape
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Example Water is not linear!
http//chemistry.tutorvista.com/
Atoms in a molecule try to spread out from one
another as much as possible to reduce the charge
repulsion between their outer electrons.
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methane, CH4
Is this the farthest that the hydrogens can get
away from each other?
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.
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Molecular Shape
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• Molecules adopt a geometry (shape) that minimizes
  e e repulsions. This occurs when e- pairs are
  as far apart as possible.

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Sample problem molecular geometry
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What is the shape of the following molecules?
commons.wikimedia.org
tetrahedral
http//winter.group.shef.ac.uk/
trigonal planar
en.wikipedia.org 
www.chriscrews.com
Bent or angular
Trigonal pyramid
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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) 3-dimensional
  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. NF32. 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 nonpolarAsymmetric
molecules with polar bonds are polar
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Naming Binary Molecules

• Write the name of the first element.
• Change the nonmetals ending to ide.
• 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

• Determine the mass of each element in the sample.
• 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 (._ _ _
  )!
• Divide the of moles of each element by the
  smallest of moles. This is the mole ratio for
  each element in the compound.
• If your answers to step 3 are whole numbers,
  these are written as the subscripts.
• 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_ 3.125 moles O
• 16.0 g

Label Properly!!
Round to thousandths (._ _ _)
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Step 3 Divide by the smallest number of moles
to obtain a mole ratio.
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 compounds
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
• 45.9 g P __1 mol_ 1.481 moles P
• 31.0 g

Label Properly!!
Round to thousandths place (._ _ _ )
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Step 3 Divide by the smallest number of
moles to obtain a mole ratio.
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, its 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
Mass of empirical formula
molecular formula
Mass of molecular formula
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Ex. The molar mass of a molecular formula is
283.88 g/mole and its empirical formula is P2O5.
Determine the molecular formula.

• Draw your chart and fill in the info from the
  problem.

use PT to calculate
from word 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.