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MIDYEAR PROJECT WORKSHEET
Lewis Structures & Molecular Geometry
Polarity and IMF
Kelly Franks, Lauren Cunningham, Juliet Hurvich
LEWIS STRUCTURES:
Definitions:
Valence Electrons: the outer level of electrons
Resonance Structure: one of two or more Lewis structures for a single molecule can be
drawn to represent a molecule
Bond Energies: Single bond < Double bond < Triple bond
Lewis Dot Diagrams
➔ an electron­configuration notation with only the valence electrons of an element are
shown, indicated by dots placed around the element’s symbol.
➔ the inner core electrons are not shown.
➔ The pair of dots representing a shared pair of electrons in a covalent bond is often
replaced by a long dash.
Writing Lewis Dot Diagrams
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.
Formal Charge and Lewis Structures
1. For neutral molecules, a Lewis structure in which there are no formal charges is
preferable to one in which formal charges are present.
2. Lewis structures with large formal charges are less plausible than those with small
formal charges.
3. Among Lewis structures having similar distributions of formal charges, the most
plausible structure is the one in which negative formal charges are placed on the
more electronegative atoms.
The Octet Rule
➔ Eight electrons in the valence shell (filling s and p orbitals) make an atom STABLE
➔ s2p6
➔ Bond formation follows the octet rule: Chemical compounds tend to form so that
each atom:
◆ by gaining, losing, or sharing electrons, has an octet of electrons in its
valence energy level.
Lewis Structures for Compounds
➔ The pair of dots between two symbols represents the shared pair of a covalent
bond.
➔ Each fluorine atom is surrounded by three pairs of electrons that are not shared in
bonds.
➔ An unshared pair, also called a lone pair, is a pair of electrons that is not involved in
bonding and that belongs exclusively to one atom.
Covalent Bonds
➔ A chemical bond in which two or more electrons are shared by two atoms.
Multiple Covalent Bonds
➔ double covalent bond or double bond:
◆ covalent bond in which two pairs of electrons are shared between two atoms
◆ shown by two side­by­side pairs of dots or by two parallel dashes
➔ triple covalent bond or triple bond :
◆ covalent bond in which three pairs of electrons are shared between two
atoms
Questions:
1. Write the lewis structure for CO2
2. Write the lewis structure for H2O
MOLECULAR GEOMETRY:
Definitions:
VSEPR theory (Valence Shell Electron Pair Repulsion): used to predict the shape of
individual molecules based upon the extent of electron­pair electrostatic repulsion
Species
Type
Model
Number of Number of Shape
Surrounding Lone Pairs
Atoms
Bond Angle
AX2
2
0
linear
180 degrees
AX3
3
0
trigonal
planar
120 degrees
AX2E1
2
1
bent
less than
120 degrees
AX4
4
0
tetrahedral
109.5
degrees
AX3E1
3
1
trigonal
pyramidal
107 degrees
AX2E2
2
2
bent
104.5
degrees
AX5
5
0
trigonal
bipyramidal
120 and 90
degrees
AX4E1
4
1
see­saw
less than 90
and less
than 120
degrees
AX3E2
3
2
t­shaped
90 and 180
degrees
AX2E3
2
3
linear
180 degrees
AX6
6
0
octahedral
90 degrees
AX5E1
5
1
square
pyramidal
less than 90
degrees
AX4E2
4
2
square
planar
90 degrees
AX3E3
3
3
t­shaped
90 and 180
degrees
AX2E4
2
4
linear
180 degrees
Questions:
1) Identify the name of the shape and bond angles of perchlorate.
2) Identify the name of the shape and the bond angles of phosphorus fluoride.
POLARITY:
Definitions:
Polarity: how equal bonding electrons are shared between elements
Polar bond or polar covalent bond: covalent bond with greater electron density around one
of the two atoms
Electronegativity: the ability to attract electrons
Determining Type of Bond
● Find the electronegativity difference between the elements
● if the difference is from 0 to 0.3, the bond is NONPOLAR COVALENT
○ these bonds share electrons
● if the difference is from 0.4 to 1.7, the bond is POLAR COVALENT
○ these bonds partially transfer electrons
● if the difference is greater than 1.7, the bond is IONIC
○ these bonds transfer electrons
Properties of Covalent Compounds
● usually soft and squishy
● not soluble in water
● does not conduct electricity
● low melting points
● low boiling points
Properties of Ionic Compounds
● combination of ions (cations/anions)
● hard and brittle
● tightly packed solids in a crystal lattice
● usually soluble in water
● conducts electricity when dissolved
● high melting points
○ Ionic bonds are very tightly bound (positive/negative attraction)
○ A LOT of energy is needed to break the bonds
Questions:
1) Classify NaCl as ionic, nonpolar covalent, or polar covalent.
2) Classify PCl3 as ionic, nonpolar covalent, or polar covalent.
INTERMOLECULAR FORCES:
Definitions:
Intermolecular Forces: attractive forces between molecules; hold atoms together
Diatomic Molecule: contains only 2 atoms
Ex­ H2, HCl, CO
Polyatomic Molecule: contains more than 2 atoms
Ex­ H2O, NH3, CH4
Ion: group of atoms with a net positive or negative charge
Cation: ion with a positive charge
if a neutral atom loses one or more electrons it becomes a cation
Anion: ion with a negative charge
if a neutral atom gains an electron it becomes an anion
Monotomic ion: contains only one atom
Ex­ Na+, Cl­
Polyatomic Ion: contains more than one atom
Ex­OH­, NH4+
Types of Intermolecular Forces
● Dipole­Induced Forces (Dispersion)
○ the attractive interaction between a polar molecule and the induced dipole
○ an induced dipole is an atom or nonpolar molecule where the separation of
positive and negative charges in the atom is due to the proximity of an ion or a
polar molecule
● London Dispersion Forces
○ attractive forces between 2 nonpolar molecules
○ very brief attraction
● Hydrogen Bonding
○
○
●
●
special type of dipole­dipole interaction
hydrogen of one molecule attracted to the directly connected oxygen, nitrogen, or
fluorine of another
○ large amount of energy for a dipole­dipole interaction; have a powerful effect on
compounds
Ion­Dipole Forces
○ attract an ion and a polar molecule to each other
○ strength depends on the size and charge of ion
○ need to be 2 different types of molecules
Dipole­dipole forces
○ attractive forces between polar molecules
○ molecules that contain dipole moments
○ the larger the dipole moment, the stronger the force
○ tend to align with opposite polarities beside one another
Forces Weakest to strongest:
1. London dispersion
2. dipole­induced
3. dipole­dipole
4. hydrogen bonding
*If you have a strong bond, you automatically have all the bonds below it*
Questions:
1. Which of these compounds is capable of forming a hydrogen bond?
MgO HF HCl NaF
2. What is the strongest intermolecular force in the molecule HCl?
3. What type of force is acted on CH4?
ANSWER PAGE
Lewis structures:
Molecular Geometry:
1. Tetrahedral; 109.5 degrees
2. trigonal bipyramidal; 120 and 90 degrees
Polarity:
1. ionic
2. polar covalent
Intermolecular Forces:
1. HF
2. dipole­dipole
3. London Dispersion