General Chemistry Topic, Senior high STEM

1. INTERMOLECULAR FORCES
AND
LIQUIDS AND SOLIDS
G R O U P 4

2. T O P I C S
:1. Kinetic molecular model of liquids and solids
2. Intermolecular forces
3. Dipole-dipole forces
4. Ion-dipole forces
5. Dispersion forces
6. Hydrogen bonding.
7. Properties of liquids and IMF
8. Surface tension
9. Viscosity
10. Vapor pressure, boiling point
11. Molar heat of vaporization
12. Structure and properties of water
13. Types and properties of solids
14. Crystalline and amorphous solids
15. Types of crystals – ionic, covalent, molecular, metallic.
16. Phase Changes – phase diagrams of water and carbon dioxide

3. KINETIC MOLECULAR MODEL
OF
LIQUIDS AND SOLIDS

5. INTERMOLECULAR FORCES
• Intermolecular forces are attractive forces between
molecules.
• The forces holding molecules together are generally
called intermolecular forces.
• Intermolecular forces are particularly important in
terms how molecules interact and form biological
organisms or even life.

7. TYPES OF
INTERMOLECULAR FORCES

9. 2. ION-DIPOLE INTERACTION:
• Is the result of an electrostatic interaction between a charged ion and a molecule
that has a dipole. It is an attractive force that is commonly found in solutions,
especially ionic compounds dissolved in polar liquids. A cation can attract the
partially negative end of a neutral polar molecule, while an anion attracts the positive
end of a polar molecule. Ion-dipole attractions become stronger as the charge on
the ion increases or as the magnitude of the dipole of the polar molecule increases.
• Attractive forces between an ion and a polar molecule

10. 3. DISPERSION FORCE:
Is the weakest intermolecular force. The dispersion force is a
temporary attractive force that results when the electrons in two
adjacent atoms occupy positions that make the atoms form temporary
dipoles. This force is sometimes called an induced dipole-induced
dipole attraction. dispersion forces are the attractive forces that cause
nonpolar substances to condense to liquids and to freeze into solids
when the temperature is lowered sufficiently.

11. 4.

12. PROPERTIES OF LIQUIDS
AND
IMF (Intermolecular forces)

13. Liquid particles have more space between them, so they are not fixed in
position. The attraction between the particles in a liquid keeps the volume of
the liquid constant.
The movement of the particles causes the liquid to be variable in shape.
Liquids will flow and fill the lowest portion of a container, taking on the shape
of the container but not changing in volume. The limited amount of space
between particles means that liquids have only very limited compressibility.

14. Surface tension: Is the energy, or work, required to increase
the surface area of a liquid due to intermolecular forces.
Since these intermolecular forces vary depending on the
nature of the liquid (e.G. Water vs. Gasoline) or solutes in
the liquid (e.G. Surfactants like detergent), each solution
exhibits differing surface tension properties. Whether you
know it or not, you already have seen surface tension at
work. Whenever you fill a glass of water too far, you may
notice afterward that the level of the water in the glass is
actually higher than the height of the glass. You may have
also noticed that the water that you spilled has formed into
pools that rise up off the counter. Both of these
phenomena are due to surface tension.

15. Viscosity

16. • Informally, viscosity is the quantity that describes a
fluid's resistance to flow. Fluids resist the relative motion
of immersed objects through them as well as to the
motion of layers with differing velocities within them.
• Viscosity is a measure of the resistance of a fluid to
deformation under shear stress.
• It is commonly perceived as "thickness", or resistance to
pouring.
• Viscosity describes a fluid's internal resistance to flow
and may be thought of as a measure of fluid friction.

17. • Formally, viscosity (represented by the symbol η "eta") is the ratio of the
shearing stress (F/A) to the velocity gradient (Δvx/Δz or dvx/dz) in a
fluid.
• η =
F̅/A
Δvx/Δz or η =
F/A
dvx/dz
• The SI unit of viscosity is the pascal second [Pa s], which has no special
name.
• The pascal second is rarely used in scientific and technical writing today.
The most common unit of viscosity is
the dyne second per square centimeter[dyne s/cm2], which is given the
name poise [P] after the French physiologist Jean Poiseuille(1799–1869).

18. •Ten poise equal one pascal second [Pa s]
making the centipoise [cP]
and millipascal second [mPa s] identical.
1 Pa s = 10 P
1,000 mPa s = 10 P
1 mPa s = 0.01 P
1 mPa s = 1 cP

19. Two quantities that are called viscosity
• The quantity defined above is sometimes
called dynamic viscosity, absolute viscosity,
or simple viscosity to distinguish it from the other
quantity, but is usually just called viscosity. The other
quantity called kinematic viscosity (represented by the
Greek letter ν "nu") is the ratio of the viscosity of a fluid
to its density.
• 𝒗 =
𝜼
𝝆

20. Differences between dynamic and kinematic viscosities
• Dynamic viscosity (also known as absolute viscosity) is the measurement
of the fluid’s internal resistance to flow while kinematic viscosity refers to
the ratio of dynamic viscosity to density. Based on the expression above,
two fluids with the same dynamic viscosities can have very different
kinematic viscosities depending on density and vice versa. As a result,
grasping the physical meaning of these two material properties may not
always be so easy.
• dynamic viscosity gives you information on the force needed to make the
fluid flow at a certain rate, while kinematic viscosity tells how fast the fluid
is moving when a certain force is applied.

21. • The SI unit of kinematic viscosity is the square meter per
second [m2/s], which has no special name. This unit is so large
that it is rarely used. A more common unit of kinematic
viscosity is the square centimeter per second [cm2/s], which is
given the name stokes [St] after the Irish mathematician and
physicist George Stokes (1819–1903). One square meter per
second is equal to ten thousand stokes.
1 cm2
/s = 1 St
1 m2
/s = 10,000 cm2
/s
1 m2
/s = 10,000 St

22. VISCOELASTICITY
• When a force (F) is applied to an object, one of four things can
happen.
• 1. It could accelerate as a whole, in which case Newton's second
law of motion would apply…
• F = ma
• This term is not interesting to us right now. We've already
discussed this kind of behavior in earlier chapters. Mass (m) is
resistance to acceleration (a), which is the second derivative of
position (x). Let's move on to something new.

23. • 2. It could flow like a fluid, which could be described by this
relationship…
• F = −bv
• This is the simplified model where drag is directly proportional
to speed (v), the first derivative of position (x). We used this in
terminal velocity problems just because it gave differential
equations that were easy to solve. We also used it in the damped
harmonic oscillator, again because it gave differential equations
that were easy to solve (relatively easy, anyway). The
proportionality constant (b) is often called the damping factor.

24. •3. It could deform like a solid according to
Hooke's law…
•F = −kx
•The proportionality constant (k) is the
spring constant. Position (x) is not the
part of any derivative nor is it raised to
any power.

25. •It could get stuck…
•F = −f
•That symbol f makes it look like we're
discussing static friction. In fluids (non-
newtonian fluids, to be specific) a term
like this is associated with yield stress.
Position (x) is not involved in any way.

26. VAPOUR PRESSURE,
BOILING POINT

27. • Vapor pressure is a liquid property related to
evaporation. In the liquid (or any substance) the
molecules have a distribution of kinetic energies related
to the temperature of the system. Because this is a
distribution there will always be a few molecules that
have enough kinetic energy to over come the attractive
potential energy of the other molecules (the
intermolecular force), and escape the liquid into the gas
phase. In an open container, these molecules will
wander off (diffuse) into the room and out into the
atmosphere. Eventually all the liquid will evaporate.
• As the temperature of a liquid increases, the vapor
pressure of the liquid increases until it equals the
external pressure, or the atmospheric pressure in the
case of an open container. Bubbles of vapor begin to
form throughout the liquid, and the liquid begins to
boil. The temperature at which a liquid boils at exactly 1
atm pressure is the normal boiling point of the liquid.
For water, the normal boiling point is exactly 100°C.
VAPOR PRESSURE, BOILING POINT

28. MOLAR HEAT OF VAPORIZATION

29. Molar heat of vaporization:
• Is the amount of heat energy required to convert
one mole of a liquid substance at its boiling point
to its gaseous state.
• Note the two important factors:
1. It's 1.00 mole of a substance
2. There is no temperature change

30. • Every substance has its own molar heat of
vaporization.
• The units for the molar heat of vaporization are
kilojoules per mole (kJ/mol). Sometimes the unit
J/g is used. In that case, it is referred to as the heat
of vaporization, the term 'molar' being eliminated.
• The molar heat of vaporization for water is 40.7
kJ/mol. To get the heat of vaporization, you simply
divide the molar heat by 18.015 g/mol.

31. • The molar heat of vaporization equation looks like this:
• q = ΔHvap mass/molar mass
• The meanings are as follows:
1. q is the total amount of heat involved
2. ΔHvap is the symbol for the molar heat of vaporization. This value is a
constant for a given substance.
3. (mass/molar mass) is the division to get the number of moles of
substance

32. 1). 49.5 g of H2O is being boiled at its boiling point of
100 °C. How many kJ is required?
• SOLUTION:
plug the appropriate values into the molar heat equation
shown above
q = (40.7 kJ / mol) (49.5 g / 18.0 g/mol)
E X A M P L E S :

33. • 2) 80.1 g of H2O exists as a gas at 100 °C. How many kJ must
be removed to turn the water into liquid at 100 °C
• SOLUTION:
• note that the water is being condensed. The molar heat of
vaporization value is used at the solid-liquid phase change,
REGARDLESS of the direction (boiling or condensing).
• q = (40.7 kJ/mol) (80.1 g / 18.0 g/mol)

34. • 3 Using the heat of vaporization for water in J/g,
calculate the energy needed to boil 50.0 g of water at its
boiling point of 100 °C.
• SOLUTION:
• multiply the heat of vaporization (expressed in J/g) by the
mass of the water involved.
• (2259 J/g) (50.0 g) = 112950 J = 113 kJ

35. The Structure and
Properties of Water

36. The Properties of Water
• Water is the most abundant compound on Earth’s
surface. In nature, water exists in the liquid, solid, and
gaseous states. It is in dynamic equilibrium between the
liquid and gas states at 0 degrees Celsius and 1 atm of
pressure. At room temperature (approximately 25
degrees Celsius), it is a tasteless, odorless, and colorless
liquid. Many substances dissolve in water, and it is
commonly referred to as the universal solvent.

38. The Phases of Water
•Similar to many other substances,
water can take numerous forms. Its
liquid phase, the most common phase
of water on Earth, is the form that is
generally meant by the word “water.”

39. Solid Phase (Ice):
• The solid phase of water is known as ice and
commonly takes the structure of hard, amalgamated
crystals, such as ice cubes, or of loosely accumulated
granular crystals, such as snow. Unlike most other
substances, water’s solid form (ice) is less dense than
its liquid form, as a result of the nature of its
hexagonal packing within its crystalline structure.
This lattice contains more space than when the
molecules are in the liquid state.

40. Liquid Phase (Water)
• Water is primarily a liquid under standard conditions (25 degrees Celsius
and 1 atm of pressure). This characteristic could not be predicted by its
relationship to other, gaseous hydrides of the oxygen family in the periodic
table, such as hydrogen sulfide. The elements surrounding oxygen in the
periodic table – nitrogen, fluorine, phosphorus, sulfur, and chlorine – all
combine with hydrogen to produce gases under standard conditions. Water
forms a liquid instead of a gas because oxygen is more electronegative than
the surrounding elements, with the exception of fluorine. Oxygen attracts
electrons much more strongly than does hydrogen, resulting in a partial
positive charge on the hydrogen atoms and a partial negative charge on the
oxygen atom. The presence of such a charge on each of these atoms gives a
water molecule a net dipole moment.

41. Gas Phase (Water Vapor)
• The gaseous phase of water is known as water vapor
(or steam) and is characterized by a transparent
cloud. Water also exists in a rare fourth state called
supercritical fluid, which occurs only in extremely
uninhabitable conditions. When water achieves a
specific critical temperature and a specific critical
pressure (647 K and 22.064 MPa), the liquid and gas
phases merge into one homogeneous fluid phase that
shares properties of both gas and liquid.

42. Phase Diagram of Water
• Water freezes to form ice, ice thaws to form liquid water,
and both water and ice can transform into the vapor state.
Phase diagrams help describe how water changes states
depending on the pressure and temperature.

43. PHASE DIAGRAM OF WATER
• The three phases of water – liquid, solid, and vapor – are shown in
temperature-pressure space.
• Note the following key points on a phase diagram:
• The critical point (CP), above which only supercritical fluids exist.
• The triple point (TP), a well-defined coordinate where the curves intersect, at
which the three states of matter (solid, liquid, gas) exist at equilibrium with
each other.
• Well-defined boundaries between solid and liquid, solid and gas, and liquid
and gas. During the phase transition between two phases (i.e, along these
boundaries), the phases are in equilibrium with each other.

44. THE POLARITY OF WATER
• The polar nature of water is a particularly important
feature that contributes to the uniqueness of this
substance. The water molecule forms an angle with an
oxygen atom at the vertex and hydrogen atoms at the
tips. Because oxygen has a higher electronegativity
than hydrogen, the side of the molecule with the
oxygen atom has a partial negative charge. An object
with such a charge difference is called a dipole
(meaning “two poles”).

45. TYPES AND PROPERTIES
OF SOLIDS

46. Solid
•A solid is a collection
of atoms or molecules that are held
together so that, under constant
conditions, they maintain a defined shape
and size. Solids, of course, are not
necessarily permanent.

47. TYPES OF SOLIDS
•Crystalline solids
• Molecular solids
• Network solids
• Ionic solids
• Metallic solids
•Amorphous solids

48. CRYSTALLINE SOLIDS
• Crystal structure determines a lot more about
a solid than simply how it breaks. Structure is
directly related to a number of important properties,
including, for example, conductivity and density,
among others. To explain these relationships, we
first need to introduce the four main types
of crystalline solids – molecular, network, ionic, and
metallic.

49. MOLECULAR SOLIDS
• Individual molecules are composed of atoms held together by
strong covalent bonds (see our Chemical Bonding module for
more about covalent bonding). To form molecular solids, these
molecules are then arranged in a specific pattern and held
together by relatively weak intermolecular forces. Examples
include ice (H2O(s) – s here stands for "solid") and
table sugar (sucrose, C12H22O11). The individual water and sugar
molecules each exist as their own independent entities that
interact with their neighbors in specific ways to create an
ordered crystalline solid.

50. NETWORK SOLIDS
• In network solids, on the other hand, there are no individually
defined molecules. A continuous networkof covalent bonds holds
together all the atoms. For example, carbon can form two different
network solids: diamond and graphite. These materials are made
up of only carbon atoms that are arranged in two different ways.
Diamond is a three-dimensional crystal that is the hardest known
natural material in the world. In contrast, graphite is a two-
dimensional network solid.
• This ability of a single element to form multiple solids is
called allotropy.

51. IONIC SOLIDS
•Ionic solids are similar
to network solids in one way: There are no
distinct molecules. But instead
of atoms held together by covalent bonds,
ionic solids are composed of positively
and negatively charged ions held together
by ionic bonds.

52. METALLIC SOLIDS
•Finally, metallic solids are a type all their
own. Although we are discussing them
last here, about three quarters of the
known elements are metals. You can read
more about these metallic elements
in The Periodic Table of
Elements module.

53. AMORPHOUS SOLIDS
• Amorphous solids are often formed when atoms
and molecules are frozen in place before they have a
chance to reach the crystalline arrangement, which would
otherwise be the preferred structure because it is
energetically favored.
• In addition, amorphous solids break unpredictably and
produce fragments with irregular, often curved surfaces,
while crystalline solids break along specific planes and at
specific angles defined by the crystal’s geometry.

54. PROPERTIES OF SOLIDS
• Crystalline solids can vary in their atomic
compositions, bonding, and structure. Together, these
attributes determine how the different solids behave under
different conditions. Solids have many different
properties, including conductivity, malleability, density,
hardness, and optical transmission, to name a few. We will
discuss just a handful of these properties to illustrate
some of the ways that atomic and molecular structure
drives function.

55. PROPERTIES OF SOLIDS
•Electrical conductivity
•thermal conductivity
•Malleability and ductility
•Melting point
•Solubility

56. ELECTRICAL CONDUCTIVITY
• Computer uses to get the electrical power it needs to run.
Those wires are made of metal, probably copper, because
metals generally have good electrical conductivity.
Electricity is essentially a flow of electronsfrom one place
to another, and in metallic bonds the outer electrons are
relatively free to move between adjacent atoms.
• The electrons are engaged in the covalent or ionic
bonds and therefore are not able to conduct electricity, or
do so only poorly. Materials that do not conduct electricity
are called electrical insulators.

57. THERMAL CONDUCTIVITY
• Heat, or thermal, conductivity is closely related to electrical
conductivity. Just as metals are good electrical conductors, you
probably know from experience that they’re good at
conducting heat too.
• For a solid to conduct heat, the movement of one molecule
or atom needs to be easily transferrable to its neighbor. The
non-directional nature of the metallic bond makes this type of
transfer relatively easy, so metals conduct heat well.
• Such solids would be expected to have low heat conductivity
and would be called heat insulators.

58. MALLEABILITY AND DUCTILITY
• Two additional properties, malleability and ductility, follow trends
similar to those for electrical and thermal conductivity. Malleability
describes the ability to hammer a solid into a sheet without
breaking it, and ductility refers to whether a solid can be stretched
to form a wire.
• Metallic malleability and ductility are a crucial reason that metals
are so useful. Their electrical conductivity would be much less
useful if it weren’t possible to stretch them into wires that could
then be bent and shaped at room temperature for an incredible
array of applications.

59. MELTING POINT
• Another way to deform a solid is to melt it. A solid’s melting
point depends on the strength of the interactions between its
components: Stronger interactions mean a higher melting point.
For molecular solids, melting means breaking the weak
intermolecular forces (the forces between different molecules), not
the strong covalent bonds that hold the
individual molecules together, so a compound like sugarcan be
easily melted on your stovetop. For network solids (held together by
covalent bonds), ionic solids(held together by ionic bonds),
and metallic solids (held together by metallic bonds), though, the
melting temperature depends on the strength of the specific bonds
in each solid.

60. SOLUBILITY
• Melting is one way of changing a solid’s shape. Another approach is
dissolving the solid into some type of liquid, in this case referred to as
a solvent. The extent to which a solid dissolves in a particular solvent
is called its solubility. Solids can be dissolved into a variety of types of
solvents, but for now we will focus on solubility in water.
• Dissolving a solid requires breaking different types of bonds for
different types of solids. Dissolving a metal requires
breaking metallic bonds, and dissolving a network solid requires
breaking covalent bonds. Both of these types of bonds are very strong
and hard to break. Therefore, metals and network solids are generally
not soluble in water.

61. TYPES OF CRYSTALS

62. CRYSTAL:
• A crystal is any solid material with its atoms, or smallest
units of matter, organized in a repeating pattern. Look at
this diagram to see an example of how atoms are
arranged in a crystal.
• There are four types of crystals:
• covalent
• Ionic
• Metallic
• molecular

63. COVALENT CRYSTALS
•Covalent crystals are crystals whose atoms are
connected with covalent bonds.
•Covalent bonds exist where the atoms share
electrons. These bonds are extremely strong and
very hard to break. Because of this, the crystals
themselves are also very strong and have high
melting points.

64. Ionic crystals
•Ionic crystals are crystals whose atoms are held
together with ionic bonds, or charged bonds. With
these ionic bonds, one atom is negatively charged
and is attracted to other atoms in the crystal that
are positively charged. They are arranged in a
pattern based on the charges. These crystals are
typically solid with a high melting point. An
example of an ionic crystal is table salt.

65. METALLIC CRYSTAL
•Metallic crystals consist of metal cations
surrounded by a "sea" of mobile valence
electrons (see figure below). These electrons,
also referred to as delocalized electrons, do not
belong to any one atom, but are capable of
moving through the entire crystal. As a result,
metals are good conductors of electricity.

66. MOLECULAR CRYSTALS
• Molecular crystals typically consist of molecules at the lattice points
of the crystal, held together by relatively weak intermolecular forces
The intermolecular forces may be dispersion forces in the case of
nonpolar crystals, or dipole-dipole forces in the case of polar crystals.
Some molecular crystals, such as ice, have molecules held together by
hydrogen bonds. When one of the noble gases is cooled and
solidified, the lattice points are individual atoms rather than
molecules. In all cases, the intermolecular forces holding the particles
together are far weaker than either ionic or covalent bonds. As a result,
the melting and boiling points of molecular crystals are much lower.
Lacking ions or free electrons, molecular crystals are poor electrical
conductors.

67. PHASES CHANGES: PHASE
DIAGRAMS OF WATER AND
CARBON DIOXIDE

68. PHASE DIAGRAMS
• A diagram showing the various phases of a system is called a phase
diagram. Phase diagrams for a pure compound such as phase
diagrams for water and carbon dioxide are phase diagrams for a
single component system. In these diagrams, pressure (P) and
temperature (T) are usually the coordinates. The phase diagrams
usually show the (P, T) conditions for stable phases.
• Phase diagrams are useful for material engineering and material
applications. With their aid, scientists and engineers understand
the behavior of a system which may contain more than one
component (compounds).

69. PHASE DIAGRAM OF WATER
• The phase diagram of water is shown here. It shows that at low
temperature, (solid) ice is the stable phase. At moderate
temperatures and high pressure, (liquid) water is the stable phase,
and at high temperature and low pressure, (gas) vapor is the stable
phase. Lines separate these phases. We discuss these items
separately below.
• The sublimation curve separate the solid from the gas. This line
indicates the vapor pressure of ice as a function of temperature.
The relationship can be shown in table or graph form. The
diagram is a sketch, and the following table gives more accurate
numbers.

70. Vapor Pressure P in mm Hg of Ice at Temperature T in K
• The vaporization curve is a plot of (equilibrium) vapor pressure P as a
function of temperature T.
T K P in mm Hg T K P in mm Hg T K P in mm Hg
190 0.00025 261 1.632 271 3.880
200 0.0012 262 1.785 272 4.217
210 0.0053 263 1.950 272.5 4.40
220 0.026 264 2.131 273 4.579
230 0.074 265 2.326
240 0.25 266 2.537
245 0.351 267 2.765
250 0.58 268 3.103
255 0.939 269 3.280
260 1.49 270 3.568

71. Vapor Pressure P in mm Hg of Water at Temperature T in K
At temperature greater than 647 K, water cannot be liquified. The fluid shares
the properties of gas. Thus, no vapor pressure beyond this temperature is
measured. The temperature of 647 K is called the critical temperature, and the
vapor pressure at this temperature is called the critical pressure.
T K P in mm Hg T K P in mm Hg T K P in mm Hg T K P in mm Hg
273 4.579* 310 47.067 370 682.07 400 682.07
274 4.926 320 79.60 371 707.27 500 19848.92
275 5.294 330 129.82 372 733.24 600 92826.40
280 7.513 340 204.96 373 760.00*
290 14.53 350 314.1 374 787.57 647 165467.20*
300 26.74 360 468.77 375 815.86

72. The melting curve or fusion
curve of ice/water is very special.
It has a negative slope due to the
fact that when ice melt, the molar
volume decreases. Ice actually
melt at lower temperature at
higher pressure. Most Canadians
skate, and the liquid formed
between the skate and ice act as a
lubricate so that the skater moves
gracefully accross the ice. The
skate apply a very high pressure
on to the ice.

73. PHASE DIAGRAM OF CARBON DIOXIDE
• The phase diagram of CO2 has some common features with that of
water: sublimation curve, vaporization curve, triple point, critical
temperature and pressure. Of course, the P and Tvalues of are unique to
carbon dioxide. The phase diagrams of water and carbon dioxide are
compared here.
• The triple point of carbon dioxide occure at a pressure of 5.2 atm (3952
torr) and 216.6 K (-56.4oC). At temperature of 197.5 K (-78.5oC), the vapor
pressure of solid carbon dioxide is 1 atm (760 torr). At this pressure, the
liquid phase is not stable, the solid simply sublimates. Thus solid carbon
dioxide is called dry ice, because it does not go through a liquid state in
its phase transition at room pressure.

74. • The critical temperature for carbon dioxide is 31.1°C, and the
critical pressure is 73 atm. Above the critical temeprature, the fluid
is called super-critical fluid.
• To be more precise, the various point of the phase diagram are
further descibed below. In the phase diagram of (a) H2O and (b)
CO2, the axes are not drawn to scale. In (a), for water, note the
triple point A (0.0098°C, 4.58 torr), the normal melting (or
freezing) point B (0°C, 1 atm), the normal boiling point C (100°C, 1
atm), and the critical point D (374.4°C, 217.7 atm). In (b), for
carbon dioxide, note the triple point X(-56.4°C, 5.11 atm), the
normal sublimation point Y(-78.5°C, 1 atm), and the critical point
Z (31.1°C, 73.0 atm).