1. The document discusses molecular shapes and intermolecular forces. It explains that molecular shape is determined by valence shell electron pair repulsion (VSEPR) theory, where electron pairs around an atom repel each other and take up positions as far apart as possible. 2. The most common molecular shapes are linear, trigonal planar, tetrahedral, trigonal pyramidal, and bent. Molecular polarity depends on bond polarity and symmetry. Polar molecules have an uneven distribution of charge while nonpolar molecules have symmetrical charge distributions. 3. Intermolecular forces include hydrogen bonding, dipole-dipole interactions, and dispersion forces. Hydrogen bonding is the strongest and occurs between molecules with hydrogen bonded to fluorine,

1. Molecular Shape
and Polarity
Intermolecular Forces of Attraction
States of Matter
VSEPR Theory

2. Atoms
•The nucleus contains the p+
and n0. It is very small but
dense and massive.
•The e- move in the energy
levels outside the nucleus
• Most atoms are more stable
when they gain, lose or share
their valence electrons in
chemical bonds = the octet
rule
REVIEW

3. Nuclear Symbols

4. Ionic Bonds make crystalline solids
– Form when e- are transferred from one atom to
another as they try to complete their outer
energy level.
– Metal Atom loses e-: gets a positive charge.
– Nonmetal Atom gains e-: gets a negative
charge.
– Positively and negatively charged atoms (or
groups of atoms) are known as ions.
– Think rocks, salts and minerals in earth’s
crust.

5. Covalent Bonds make molecules
• When valence e- are shared (in pairs) by atoms
instead of being transferred.
• Equal sharing of e- = nonpolar covalent
• Unequal sharing of e- = polar covalent
• The structure that results is called a molecule.
• Covalent compounds tend to be liquids and gases
at room temperature because of weak forces of
attraction between molecules.
• think water, oceans, atmosphere and carbon-
based (organic) molecules.

6. Water molecules are Polar
Polar Covalent Bonds: Unequal sharing of electrons
results in one side of the molecule being “negative”
and the other side being “positive”
 There is a greater e- density
around one atom than there is
around the other
 Which one in H2O?
 Oxygen “pulls”
e- more strongly
(higher electronegativity)
 Partial charges form!!!
Negative end
Positive end

7. (–)
O
H
H
(+) (+)
Chemists often use arrows to show the PULL of
negative charge toward one end of the molecule. A
polar molecule is called a DIPOLE.

8. Bond Type by Electronegativity
Difference
• You can estimate
the bond type by
subtracting the
difference in EN
values of the
bonding atoms.
• Look up EN values
& subtract
EX: H = 2.1, O = 3.5
∆EN = 1.4 = polar
covalent
∆EN
electronegativity
difference
Bond
Type
≤0.4 Non Polar
Covalent
Between
0.5 to 1.7
Polar
Covalent
≥1.9 Ionic

9. Ionic vs Molecular & States of
Matter
• Much of chemistry comes down to opposite
charges attract and like charges repel.
• This means that PROPERTIES like physical
state and melting and boiling points DEPEND
ON HOW STRONGLY THE BASIC UNITS ARE
ATTRACTED TO EACH OTHER!
• Solids melt and liquids boil when the attractive
forces between the particles are broken….
 Stronger attractive forces result in higher
MPts and BPts.

10. Ionic vs Molecular & States of
Matter
• The force of attraction between ions is quite
strong – ionic crystals are hard, brittle solids
with high MPts – it takes more energy to break
these bonds.
• The force of attraction between individual
molecules are called weak forces – many
molecular substances are gases or volatile
liquids at room temp. and MPts are low.

11. INTERMOLECULAR FORCES:
Weak attractive forces between molecules
• Chemists call intermolecular forces of attraction Van der
Waals forces, after the scientist who discovered them.
• IM forces vary in strength and change the properties of
molecular substances.
• IM forces are due to polarity in molecules – differences
depend on whether dipoles are permanent or temporary.
• IM forces allow molecules to come together and form
liquids, like water, and a few covalent solids.

12. States of Matter
• Retains a fixed
shape, rigid.
• Particles locked
into place.
• Not
compressible.
• Dense-little free
space between
particles
• Take shape of
container
• Particles are
fluid, can slide
past each other
• Not compressible
• Dense-little free
space between
particles
• Expand to fill
their container
• Particles are
fluid, move rapidly
& randomly
• VERY
compressible.
• Lots of free space
between particles.

13. Hydrogen Bonds Exist Between
Water Molecules
 Form between (-) oxygen
atom of a polar molecule and
(+) hydrogen atom of a
neighboring molecule.
 The attractions between
polar water molecules are so
significant we call them
hydrogen bonds even though
they are not as strong as
true covalent or ionic bonds.
 Without hydrogen bonds
water would never be a liquid
O
H

14. Interaction Between Water
Molecules
Negative Oxygen end of one water molecule is
attracted to the Positive Hydrogen end of another
water molecule to form a HYDROGEN BOND

15. Hydrogen bond
 Hydrogen
bonds are the
strongest type
of
intermolecular
forces
 intermolecular
forces act
BETWEEN the
molecules.
These are
covalent bonds

16. White = H and Red = O
Hydrogen Bonds –
Why we are here…
• Hydrogen bonds make water a liquid at room temperature
• Hydrogen bonds allow DNA to unzip during replication
•

17. Video: Geckos
• https://youtu.be/YeSuQm7KfaE

18. Video: Van der Waals Forces
• https://youtu.be/3yXrHlLZ4LI

19. Types of Intermolecular Forces
from strongest to weakest
1. Hydrogen Bonding: H with F O N
Attraction between two molecules containing one
of the very polar bonds of O-H, N-H & H-F
2. Dipole-Dipole forces:
Attraction between two polar molecules.
3. London Dispersion Forces:
Attraction between two non-polar molecules due
to temporary, induced dipoles.

20. Hydrogen Bonds
Hydrogen bonds, especially strong dipole−dipole
attractions, occur between
• polar molecules containing hydrogen atoms bonded to
very electronegative atoms such as fluorine (F), nitrogen
(N), and oxygen (O).
• a hydrogen atom with a partial positive charge attached
to a N, O, or F with a partial negative charge.

21. Dipole−Dipole Attractions
Polar molecules are attracted to each other by
dipole−dipole attractions when the positive end of one
dipole is attracted to the negative end of a second
dipole, such as the attractive forces between two
molecules of H—Cl.

22. Dispersion Forces
Dispersion forces, very weak attractive forces that
occur between nonpolar molecules,
• occur when movement induces a temporary
distortion of the electrons in a molecule, creating a
temporary dipole.
• make it possible for nonpolar molecules to exist as
liquids and solids.

23. Bonding and Attractive Forces

24. Learning Check
Indicate which major type of molecular interaction,
dipole−dipole attractions, hydrogen bonds, or
dispersion forces, is expected between each of the
following.
A. NF3
B. Cl2
C. HF

25. Solution
Indicate which major type of molecular interaction,
dipole−dipole attractions, hydrogen bonds, or
dispersion forces, is expected between each of the
following.
A. NF3 dipole−dipole attractions
B. Cl2 dispersion forces
C. HF hydrogen bonds

26. A few elements make most molecules
H
H Be B C N O F
P S Cl
STOP AND DRAW MOLECULES TOGETHER

27. VSEPR Theory
• Valence Shell Electron Pair Repulsion
Theory
• States that:
Repulsion between sets of valence electrons
surrounding an atom causes them to be
oriented as far apart as possible.
VSEPR theory postulates that the lone pairs occupy
space around the central atom just like bonding
pairs, but they repel other electron pairs more
strongly than bonding pairs do.

28. The water molecule has two unshared
electron pairs

29. Lewis Structures
• The pair of dots representing
the shared pair of electrons in
a covalent bond is often replaced
by a long dash.
• 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.
Shared pair
(covalent bond) Lone
pair
•
•

30. The five most common shapes of
small molecules
1. Linear
2. Trigonal Planar
3. Tetrahedral
4. Trigonal Pyramidal
5. Bent

31. Molecular Geometry
• The shape of molecules affects the
physical & chemical properties of
molecular (covalent) substances.
• All molecules have a symmetrical shape
because the bonds & atoms are
arranged with equal distances
separating atoms that are NOT bonded
to each other.
• WHY? A repulsive force exists
between e- pairs in molecules.

32. How do molecules get their shapes?
• Valence e- surrounding an atom may
be shared in pairs OR left
unshared.
• Both bonded and unbonded e- pairs
will repel each other – LIKE
CHARGES REPEL.
• Unshared e- pairs repel each other
VERY STRONGLY!!!

33. Hybrid Orbitals – form when s & p orbitals mix, creating
symmetry around the central atom in a molecule

34. Molecular Geometry
Common Molecular Shapes
Linear: The atoms of a linear molecule are connected in a
straight line.
All 2 atom molecules are linear (O2, HCl).
Many 3 atom molecules are also linear. (CO2)
6-26

35. Molecular Geometry
• Trigonal Planar: Molecules
have a triangular, flat shape.
*4 atom molecules
EX: BCl3
A central atom
bonded to 3
other atoms.
No unshared e- pairs on central atom!

36. Molecular Geometry
• Tetrahedral: A shape that
has four surfaces.
*5 atom molecules.
Tripod shaped.
Ex. CH4 (Methane)
All four sides are identical.
No unshared e- pairs on
central atom!

37. Molecular Geometry
• Trigonal Pyramidal: The
molecule has a central atom
that is bonded to three other
atoms and has one unshared
pair on the central atom.
*4 atom molecules
Ex. NH3 (Ammonia)

38. Molecular Geometry
• Bent: 2 unshared e-
pairs on the central
atom result in a
slightly smaller bond
angle, due to an even
greater repulsion
force.
*3 atom molecules
Ex. H2O

39. Molecular Polarity
• Depends on both bond
polarity and molecular
geometry.
 If all bonds are
non-polar, the molecule
is always non-polar.
 If bonds are polar, but there is symmetry in the molecule
so that the polarity of the bonds cancels out, then the
molecule is non-polar. (Ex: CO2, CCl4)
 If bonds are polar but there is no symmetry such that
they DO NOT cancel each other out, the overall molecule
is polar. (Ex: H20, CH3Cl)
Chapter 6 – Section 5: Molecular Geometry

40. Polarity of Molecules—Nonpolar
In a nonpolar molecule, all the bonds are nonpolar,
H2, Cl2 and CH4 are nonpolar because they contain only
nonpolar bonds.

41. Polarity of Molecules—Nonpolar
A nonpolar molecule also occurs when polar bonds
(dipoles) cancel each other because of a symmetrical
arrangement.
Molecules such as CO2 and CCl4 contain polar bonds with
dipoles that cancel each other out.

42. Polarity of Molecules—Polar
A polar molecule occurs when the dipoles from
individual bonds do not cancel each other out.
For molecules with two or more electron groups, the
shape (such as bent or trigonal pyramidal) determines
whether or not the dipoles cancel.

43. Polarity of Molecules—Polar
Examples of polar molecules include HCl, H2O, and NH3.
• HCl is linear and contains a polar bond.
• H2O is bent and contains two polar bonds as well as two
lone pairs on oxygen.

44. Polarity of Molecules—Polar
• NH3 is trigonal pyramidal and contains three polar
bonds and a lone pair on nitrogen.

45. Polarity of Molecules—Polar
• CH3F is tetrahedral and contains three nonpolar
bonds and a polar bond.

46. White = H and Red = O
Water
• A water molecule (H2O), is made up of three
atoms --- one oxygen and two hydrogen.
• Each H atom forms a single covalent bond
with the O atom.
H
H
O
H
H