K-12 Curriculum General Chemistry 2 Senior High School STEM 11/12

1. GENERAL CHEMISTRY II

2. General Chemistry 2 – Senior High School (STEM)
EQ: Why do solids and liquids behave differently?

3.  The Kinetic Molecular Theory explains the
properties of solids and liquids in terms of
intermolecular forces of attraction and the
kinetic energy of the individual particles.

4. 1. All matter is made up of tiny particles.
2. These are particles are in constant motion.
3. The speed of particle is proportional to
temperature. Increased temperature means
greater speed.
4. Solids, liquids, and gases differ in distances
between particles, in the freedom of motion of
particles, and in the extent to which the
particles interact.

6. a. Compare the distances among molecules in the gas,
liquid and solid and rank the phases in increasing
distance between particles.
b. Describe the characteristic movement of the particles
of gas, liquid and solid.
c. How are the molecules of gas, liquid and solid
arranged?
d. Arrange the three phases of matter in order of
increasing volume of empty space between its
molecules.

10. General Chemistry 2 – Senior High School (STEM)
EQ: How is intermolecular forces defined by nature
of particles?

11.  INTERMOLECULAR FORCES are attractive
forces between molecules or particles in the
solid or liquid states.

12.  INTERMOLECULAR FORCES (IMF) are
relatively weaker than the forces within the
molecules forming bonds (intramolecular
forces)
 Intramolecular Forces hold atoms together in
a molecule.

13.  The intermolecular forces of attraction in a
pure substance are collectively known as van
derWaals forces.
1. Dipole-dipole
2. Hydrogen bonding
3. Ion-dipole
4. London dispersion
5. Dipole-induced dipole force

14.  Dipole-dipole forces exist between polar
molecules. One end of a dipole attracts the
oppositely charged end of the other dipole.

15.  It is a special and very strong type of dipole –
dipole force that exists between a hydrogen
atom bound to a small and highly
electronegative non-metal atom.
 Hydrogen bond occurs in polar molecules
containing H and any of highly
electronegative elements, in particular
Nitrogen, Fluorine, and Oxygen.

17.  It acts between an ion
(either positive or negative)
and a polar molecule.
 This explains the solubility of
ionic compounds in water,
which is polar molecule.

18.  The ions and the oppositely charged ends
of the polar water molecules overcome the
attraction between ions themselves.
Each ion becomes separated and water
molecules cluster around it.

19.  It is the weakest type of intermolecular force.
 When two non-polar molecules approach
each other, an instantaneous dipole moment
forms.
 This force is sometimes called an induced
dipole-induced dipole attraction.

21.  Interaction between Polar and non-polar
molecules.

22. What type of intermolecular force will act in the
following substances? Justify your answer.
1. sulfur dioxide (SO2)
2. nitrogen gas (N2)
3. hydrogen fluoride (HF)
4. carbon dioxide (CO2)
5. neon gas (Ne)
6. magnesium chloride (MgCl2) dissolved in water
(H2O)

23. General Chemistry 2 – Senior High School (STEM)
EQ: How do intermolecular forces influence the
properties of liquids?

24.  Liquids do not have a
simple or regular
structure, but many
of their properties
can be explained
qualitatively by
viewing them at the
particulate level.

25. General
Properties
of Liquids
Surface
Tension
Viscosity
Vapour
Pressure
Boiling Point
Heat of
Vaporization
Capillary
Action

26.  It is the measure of the elastic force in the
surface of a liquid.
 It is the amount of energy required to stretch
or increase the surface of a liquid by a unit
area.
 It is manifested as some sort of skin on the
surface of a liquid or in a drop of liquid.

27.  Surface tension allows needles and paper
clips to float in water if placed carefully on the
surface. It also explains why drop of water are
spherical in shaped

29.  These intermolecular forces tend to pull the
molecules into the liquid and cause the
surface to tighten like an elastic film or “skin”.

30.  Molecules within a liquid
are pulled in all directions
by intermolecular forces.
 Molecules at the surface
are pulled downward and
sideways by other
molecules, not upward
away from the surface

31.  The liquids that have strong Intermolecular
forces also have high surface tension.

32.  Capillary action is the tendency of a liquid to
rise in narrow tubes or be drawn into small
openings such as those between grains of a
rock.
 Capillary action, also known as capillarity, is
a result of intermolecular attraction between
the liquid and solid materials.

33.  Capillary action is shown by water rising spontaneously in
capillary tubes. A thin film of water adheres to the wall of the
glass tube as water molecules are attracted to atoms making
up the glass (SiO2).

36.  Two types of forces are involved in capillary action:
 Cohesion is the intermolecular attraction between
like molecules (the liquid molecules).
 Adhesion is an attraction between unlike molecules
(such as those in water and in the particles that
make up the glass tube).
 These forces also define the shape of the surface of
a liquid in a cylindrical container (the meniscus!)

37.  When the cohesive forces between the liquid
molecules are greater than the adhesive forces
between the liquid and the walls of the container,
the surface of the liquid is convex.
 When the cohesive forces between the liquid
molecules are lesser than the adhesive forces
between the liquid and the walls of the container,
the surface of the liquid is concave.

39.  It is defined as the resistance
of a liquid to flow.
 It is loosely referred to as the
thickness or thinness of a
liquid.
 Syrup and oil flow more
slowly than water and are
thus described as more
viscous.

40.  The viscosity of liquid
depends on their
intermolecular attraction.
 The stronger the
intermolecular force, the
higher is the liquid’s
viscosity

41.  Long-chained substances like
oil have greater
intermolecular forces
because there are more
atoms that can attract one
another, contributing to the
substance’s total attractive
forces.

42.  Honey, a concentrated
solution of sugar, is also
highly viscous because of the
hydrogen bonding that
forms as a result of the
numerous- OH groups of
sugar molecule.

43.  It is the pressure exerted by its vapor when in
equilibrium with liquid or solid.
Example:
 When liquid or solid substance is made to
evaporate in a closed container, the gas
exerts a pressure above the liquid.

44.  Substances with relatively strong
intermolecular forces will have low vapor
pressure because the particles will have
difficulty escaping as a gas.
Example:
1. Water (H2O), (Hydrogen Bonding) has vapor
pressure of 0.03 atm.
1. Ethyl Ether (C4H10O), dipole-dipole & London
Force ) has vapor pressure at 0.68 atm.

45.  The boiling point of a liquid is the
temperature at which its vapor pressure is
equal to the external or atmospheric
pressure.
 Increasing the temperature of a liquid raises
the kinetic energy of its molecules, until such
point where the energy of the particle
movement exceeds the intermolecular forces
that hold them together.

46.  The liquid molecules then transform to gas
and are seen as bubbles that rises to the
surface of the liquids and escape to the
atmosphere.
 Then temperature at which a liquid boils
under 1 atmospheric pressure (1atm) is
referred to as its normal boiling point.

47.  At higher altitude, the atmospheric pressure
is lower, hence, the boiling point will
subsequently decrease.
 The greater intermolecular force, the higher
the energy needed to increase the kinetic
energy of the molecules to break these
forces.

49.  Molar Heat of vaporization ( Hvap) is the
amount of heat required to vaporize one
mole of substance at its boiling point.
 The application of heat disrupts the
intermolecular forces of attraction of the
liquid molecules and allows them to vaporize.

50.  Boiling point generally increases as molar heat
of vaporization increases.
 The Hvap is also determined by the strength
of intermolecular forces between molecules.

52.  At room temperature,
pure water is a
colorless, odorless and
tasteless liquid.
 It turns to ice, its solid
form at 00 C and 1 atm.
 At 1000 C, it become
gas, commonly known
as steam.

53. 1.Water is a good solvent.
2.Water has a high specific heat.
Specific heat is the amount of heat or energy
needed to raise the temperature of one gram of
a substance by 1o C.
3.The boiling point of water unusually high.

54. 4. Solid water is less dense, and in fact floats on liquid
water.
 Unlike all other liquids, the molecules in solid water
are actually farther apart than they are in liquid water.
 When solid water forms, the hydrogen bonds result
in a very open structure with unoccupied spaces,
causing the solid to occupy a larger volume than the
liquid.
 This makes ice less dense than liquid water, causing
ice to float on water.

57. General Chemistry 2 – Senior High School (STEM)
EQ: How do you describe solids?

58.  Solid can be classified as crystalline or
amorphous based on the arrangement of
their particles.
 Crystalline solids have highly regular
arrangement of particles, while amorphous
solids have considerable disorder in their
structure.

59.  Amorphous solids, such
a glass, are formed
rapidly that its
constituent particles do
not have time to align
or organize into a more
crystalline lattice.

61.  Crystalline Solids have well-defined crystal
lattice.
A lattice is a three-dimensional system of
points designating the positions of the
components (ions, atoms, or molecules) that
makeup a crystal.

63.  A unit cell is the smallest repeating unit of
lattice.

64. TYPES COMPONENTS
THAT OCCUPY
THE LATTICE
POINTS
TYPE OF
INTERACTION
BETWEEEN
COMPONENTS OF
LATTICE
TYPICAL PROPERTIES EXAMPLES
IONIC Ions Ionic Hard, high melting
point; insulating as solid
but conducting when
dissolved.
NaCl
MOLECULAR Discrete
molecules
Dipole-dipole or
London dispersion
Soft; low melting point Ice, dry ice
METALLIC Metal atoms Delocalized covalent Wide range of hardness
and melting points
Silver, Iron,
Brass
NETWORK Nonmetal
atoms
Directional covalent Hard, high melting point Diamond
GROUP 8A Noble gases London dispersion
forces
Very low melting point Argon

66. General Chemistry 2 – Senior High School (STEM)
EQ. When does equilibrium exist between the
phases of a substance?

67.  Phase Changes are transformations of matter
from one physical state to another.
 They occur when energy is added or removed
from a substance.
 They are characterized by changes in molecular
order; molecules in the solid phase have the
greatest order, while those in the gas phase have
the greatest randomness or disorder.

68.  What changes in molecular order occur
during phase changes?

70.  How does a change in energy affect phase
changes?

71.  How does a change in energy affect phase
changes?

72.  How can this effect be achieved using CO2 or
dry ice?

73. Carbon dioxide cannot exist as a liquid at
atmospheric pressure, the dry ice sublimates
and instantly produces a gas, condensing water
vapor, and creating a thick white fog.

74.  What does LPG stand
for? How can a gas be
liquefied?
 What conditions are
needed to convert a gas
into a liquid?

75.  Liquefied petroleum gas or
liquid petroleum gas (LPG
or LP gas), are flammable
mixtures of hydrocarbon
gases.
 It is used as fuel in heating
appliances, cooking
equipment, and vehicles

76.  It is a graphical representation of the physical states
of a substance under different conditions of
temperature and pressure.
 It gives the possible combinations of pressure and
temperature at which certain physical state or
states a substance would be observed.

78.  Phase diagrams are plots of pressure (usually
in atmospheres) versus temperature (usually
in degrees Celsius or Kelvin).
1. ThreeAreas (Solid, Liquid, Gas)

79.  The three areas are marked solid, liquid, and
vapor. Under a set of conditions in the
diagram, a substance can exist in a solid,
liquid, or vapor (gas) phase.

81.  The lines that serve as boundaries between
physical states represent the combinations of
pressures and temperatures at which two
phases can exist in equilibrium.
 In other words, these lines define phase
change points.

82.  The green line divides the solid and liquid
phases, and represents melting (solid to
liquid) and freezing (liquid to solid) points.

83.  Melting (or freezing) curve – the curve on a
phase diagram which represents the
transition between liquid and solid states.
 It shows the effect of pressure on the melting
point of the solid. Anywhere on this line,
there is equilibrium between the solid and the
liquid.

84.  The blue line divides the liquid and gas
phases, and represents vaporization (liquid to
gas) and condensation (gas to liquid) points.

85.  The curve on a phase diagram which
represents the transition between gaseous
and liquid states. It shows the effect of
pressure on the boiling point of the liquid.
 Anywhere along this line, there will be
equilibrium between the liquid and the vapor.

86.  The red line divides the solid and gas phases,
and represents sublimation (solid to gas) and
deposition (gas to solid) points.

87.  The curve on a phase diagram which
represents the transition between gaseous
and solid states.
 It represents the effect of increased
temperature on a solid at a very low constant
pressure, lower than the triple point.

88.  The triple point is the combination of
pressure and temperature at which all three
phases of matter are at equilibrium.
 It is the point on a phase diagram at which
the three states of matter coexist. The lines
that represent the conditions of solid-liquid,
liquid-vapor, and solid-vapor equilibrium
meet at the triple point

90.  The critical point terminates the liquid/gas
phase line. It is the set of temperature and
pressure on a phase diagram where the liquid
and gaseous phases of a substance merge
together into a single phase.
 Beyond the temperature of the critical point,
the merged single phase is known as a
supercritical fluid.

94. Constructing a Phase Diagram
Visualize a substance with the following points on the phase diagram: a triple
point at 0.05 atm and 150 K; a normal melting point at 175 K; a normal boiling
point at 350 K; and a critical point at 2.0 atm and 450 K.
The solid liquid line is “normal” (meaning positive sloping).
For this, complete the following:
1. Roughly sketch the phase diagram, using units of atmosphere and Kelvin.
Label the area 1, 2, and 3, and pointsT and C on the diagram.
2. Describe what one would see at pressures and temperatures above 2.0 atm
and 450 K.
3. Describe the phase changes from 50 K to 250 K at 1.5 atm.
4. What exists in a system that is at 1 atm and 350 K?
5.What exists in a system that is at 1 atm and 175 K?