Chemistry midterm 2: Liquids and solids Flashcards
(84 cards)
1
Q
describe a phase
A
a physically distinct, homogeneous part of a system
comprised of one physical state of matter
2
Q
What is potential energy
A
stored energy, which in the form of attractive forces draws particles together
3
Q
What is kinetic energy
A
the energy of motion which tends to disperse particles
4
Q
Describe entropy of gases
A
Gases have very high molar entropy (disorder), liquids much less, solids even less.
5
Q
Describe enthalpy
A
Attractive intermolecular forces caused the particles re-arrange themselves going from gas to liquid to solid, manifested by enthalpy changes. Stronger intermolecular forces lead to larger enthalpy changes.
6
Q
Phase changes involve
A
the forming, breaking or changing the strength of intermolecular forces
7
Q
Intermolecular forces are
A
the attractive, electrostatic forces that exist between all molecules, ions, and atoms. They are relatively weak in
comparison to bonding forces.
8
Q
Transitions between solid, liquid, and gaseous states of a substance occur
A
when conditions of temperature or pressure favor the associated changes in intermolecular forces
9
Q
Describe ion-ion forces
A
- are some of the strongest forces between molecules
- Oppositely charged ions attract each other
- non-directional; each ion interacts equally strongly with all of its neighbours
10
Q
Describe ion-dipole forces
A
- The attachment of water to solute particles is called hydration
- Hydration of ions is due to the polar character of the H2O molecule
- Note which end of the water is attracted to either an anion or cation. Remember water has a permanent
dipole - 1/r^2 so atoms need to be fairly close
11
Q
What are polar moelcules?
A
polar molecules have a partial positive charge on one side and a partial negative charge on the other side of the molecule
12
Q
What is a dipole
A
Separation of charge for polar molecules
13
Q
What is dipole-dipole attraction?
A
—the electrostatic force between the partially positive end of one polar molecule and the partially negative end of another
14
Q
Polar molecules have a permanent
A
dipole
15
Q
What determines polarity
A
Bond polarity and molecular geometry (shape)
16
Q
For dipole dipole interactions how do molecules arrgane themselves
A
Lowest energy, least repulsive configuration
17
Q
The permanent dipole adds to the attractive forces between the molecules called a dipole-dipole interaction raising the
A
Boiling point and melting points relative to nonpolar molecules of similar size and shape
18
Q
What is one of the main factors that determine which
compound will have a higher boiling point
A
The magnitude of the dipole moment. The stronger the attractions between the atoms or molecules, the more energy it will take to separate them.
19
Q
Boiling a liquid requires adding enough energy to overcome all the attractions between the particles. However, it does not require breaking the
A
Covalent bonds
20
Q
What are hydrogen bonds
A
- Very strong intermolecular forces
- stronger than dipole-dipole or dispersion forces
21
Q
Substances that can hydrogen bond will have higher
A
boiling and melting points than substances that cannot
22
Q
How does the strength of hydrogen bonds compare to chemical bonds?
A
- hydrogen bonds are not nearly as strong as
chemical bonds - 2–5% the strength of covalent bonds
23
Q
What are examples of hydrogen bonding in nature?
A
- trees are held upright by hydrogen bonds: Cellulose
molecules (which have many -OH groups) strengthen
wood. - Cell division: Hydrogen bonding, though weaker than covalent bonding, allows DNA to readily give way in cell division without affecting the covalent bonds in DNA
24
Q
Describe dipole-induced dipole interactions
A
Polar molecules with a permanent dipole can interact with a
nonpolar molecule creating an induced dipole (for example, as when oxygen dissolves in water)
- When atoms are far apart they do not influence one other.
- When atoms are close together, the instantaneous dipole in one atom induces a dipole in the other
25
What are London dispersion forces
- Dispersion forces are the interactions that arise from the motion of electrons in atoms that cause unequal electron distribution forming instantaneous dipoles.
- Dispersion forces explain why even nonpolar noble gas can form liquids
26
What is the polarizability of a particle
- the ease with which its electron cloud is distorted
- Smaller particles are less polarizable than larger ones because their electrons are held more tightly
27
What is the relationship between dispersion forces and size
- dispersion forces increase with larger and more polarizable molecules
- As the molar mass increases, the number of electrons increases. Therefore, the strength of the dispersion forces increases
- The stronger the attractive forces between the molecules, the higher the boiling point will be
28
What is the relationship between dispersion forces and shape
- The more elongated the molecule, the greater the area for interaction. Therefore, the strength of the dispersion forces
increases
29
What is the general rule for comparing intermolecular forces
For molecules of approximately equal mass and size, the strengths of the intermolecular attractions increase with increasing polarity
30
What are the three rules when comparing intermolecular forces
1. Dispersion forces are present in ALL molecules, whether they are polar or nonpolar
2. When the molecules being compared have roughly the same numbers of electrons and shape any difference in the magnitudes of the attractive forces are attributed to dipole-dipole forces (or other types of forces)
3. When the molecules differ widely in numbers of electrons the dispersion forces tend to be the decisive ones
31
What are 5 rules for comparing IF's
1. Dispersion forces are present in all substances and become greater in magnitude with increasing numbers of electrons.
2. The strength of dispersion forces also depends on molecular shape.
3. Dipole-dipole forces add to the effect of dispersion forces and are found in polar molecules
4. Hydrogen bonds also add to the effect of dispersion forces and tend to be stronger than dipole-dipole and dispersion forces
5. None of the intermolecular forces are as strong as an ionic or covalent bond
32
What is viscosity
Viscosity is liquid’s resistance to flow: the higher the viscosity of the liquid, the more sluggish the flow. The stronger the intermolecular forces and cohesive forces, the greater the viscosity
33
With increasing temperature, the viscosity
decreases, because at higher temperatures, these
interactions are weaker as the molecules move more rapidly and their kinetic energies are better able to overcome the forces that hold them together
34
What is the viscosity of pitch
Pitch has a viscosity 10^11 times that of water
35
What is surface tension
the energy required to increase the surface area of a liquid
36
The stronger the forces between the particles, the higher the
Surface tension
37
Surface molecules experience a new attraction
towards the bulk
38
Why do liquids with high surface tensions tend to "bead"
Spheres have the lowest surface area to volume area of any shape
39
Who is considered the surface science pioneer
Auntie Agnes.
Municipal High School for Girls, Brunswick
Langmuir Trough –Surface Tension Apparatus
Surface Tension (1891) – Nature 46, 437
Adhesion of Liquids on Glass (1898)
Contact Angles and the Flow of Fluids (1914)
- she developed an instrument known as a trough for measuring surface contaminants and their effects. With the support of renowned scientist Lord Rayleigh, her 1891 paper showed that surface contamination significantly reduces surface tension, and also that changing the characteristics of the surface (compressing or expanding it) also affects surface tension
40
What are cohesive forces
- Intermolecular forces between like molecules
such as within a liquid like water. The various IMFs between identical molecules of a substance are examples of cohesive forces
41
What are adhesive forces
Intermolecular forces between unlike molecules such as between a liquid and a solid surface
42
What is capillary action
- the ability of a liquid to flow against gravity and spontaneously rising in a narrow tube or porous
structure 10.2 Capillary Action
43
What is wetting
the extent to which a liquid spreads into a thin film
on a surface depends on adhesive forces
44
What is superhydrophobicity
materials that are difficult to wet with contact angles greater than 150 °C and usually characterized by surface roughness
45
When the rate of condensation becomes equal to the rate of vaporization
neither the amount of the liquid nor the amount of the vapor in the container changes. The vapor in the container is then said to be in equilibrium with the liquid. Keep in mind that this is not a static situation, as molecules are continually exchanged between the condensed and gaseous phases
46
What is dynamic equilibrium
the status of a system in which reciprocal processes (for example, vaporization and condensation) occur at equal rates
47
What are endothermic changes
- processes that require an input of energy
48
What is enthalpy of vaporization
heat added to a liquid to transform it into a
gas by breaking attractive intermolecular forces
49
What happens during evaporation
- increased temperature
- increased surface area of the liquid
- decreased strength of intermolecular forces
50
What are exothermic changes
processes that releases energy
51
What is enthalpy of formation
heat added to a solid to transform it into a liquid by reducing attractive intermolecular forces
52
Within a phase, heat flow is accompanied by
a change in temperature, since the average KE of the particle’s changes.
53
During a phase change, heat flow occurs at a
constant temperature, as the average distance between particles changes
54
In a closed flask, the system reaches a state of dynamic equilibrium, where molecules
are leaving and entering the liquid at the same rate
55
The vapor pressure is
the pressure exerted by the vapor on the liquid. The
pressure increases until equilibrium is reached; at equilibrium the pressure is constant.
56
the partial pressure of molecules in the gaseous state at a
given temp is
Vapour pressure
57
The rate of vaporization depends on
temperature, IFs and the surface area of the liquid, and is independent of the rate of condensation, which depends on the vapour pressure.
58
How do IF's affect vapour pressure
- Relatively strong intermolecular attractive forces will serve to impede vaporization as well as favoring “recapture” of gas-phase molecules when they collide with the liquid surface, resulting in a relatively low vapor pressure.
- Weak intermolecular attractions present less of a barrier to vaporization, and a reduced likelihood of gas recapture, yielding relatively high vapor pressures. The following example illustrates this dependence of vapor pressure on intermolecular attractive forces
59
How does temp affect vapour pressure
- As temperature increases, the vapor pressure of a liquid also increases due to the increased average KE of its
molecules
- at any given temperature, the molecules of a substance experience a range of kinetic energies, with a certain fraction of molecules having a sufficient energy to overcome IMF and escape the liquid (vaporize)
- At a higher temperature, a greater fraction of molecules have enough energy to escape from the
liquid
- The escape of more molecules per unit of time and the greater average speed of the molecules that escape both contribute to the higher vapor pressure
60
When the vapor pressure increases enough to equal the external atmospheric pressure, the liquid reaches its
- Boiling point.
- The boiling point of a liquid is the temperature at which its equilibrium vapor pressure is equal
to the pressure exerted on the liquid by its gaseous surroundings. For liquids in open containers, this pressure
is that due to the earth’s atmosphere
- The normal boiling point of a liquid is defined as its boiling point when surrounding pressure is equal to 1 atm
61
What is a volatile substance
has a vapor pressure at ordinary temperatures. These substances have weak IFs
62
As the external pressure on a liquid increases
the boiling point increases
63
What is sublimation
phase transition of a substance from its solid to gaseous phase
64
What is heat of sublimation
heat required to bring one mole of a substance from its solid to gaseous state
65
What is deposition
The reverse of sublimation is called deposition, a process in which gaseous substances condense directly into the
solid state, bypassing the liquid state.
66
What are phase diagrams
the stable phases that a substance occupies at varying
pressure and temperatures
67
What are boundaries
the temperatures and pressures at which two phases are in equilibrium
68
What is the triple point
the temperature and pressure at which three states are in
equilibrium [can be multiple].
69
If we change P&T on phase diagram
- we can cause the system to cross phase boundaries
70
Moving over a transition by increasing T, at constant P, results in
an enthalpy increase. These transitions are endothermic.
71
Moving over a transition by increasing P, at constant T, results in
an increase in density
72
What is a fusion curve
The Fusion Curve has a positive slope. As we increase the pressure on the liquid, the denser state is favoured which is the solid state
73
What is the critical point
liquid and vapor phases become indistinguishable and
form a supercritical fluid with properties of liquids
and gases
74
What is critical temperature
temperature above which liquid cannot
exist no matter the pressure applied
75
What is critical pressure
maximum pressure for gas to transition to liquid 10.3 Critical Point and Supercritical Fluids
76
What are the 2 orders of solids
1. Crystalline: atoms in ordered matrix
2. Amorphous: short-range order
- solids melt over range of temperatures
- properties depend on preparation
77
What are molecular solid properties
- Intermolecular Forces : Van der Waals
interactions, Dipole-Dipole and H-Bonding
- Properties include: Low Boiling and Melting Points, Electrical insulators, Soluble in solvents of similar polarity
- Boiling point depends on: number of electrons, as expected
- Melting point depends on: strength of intermolecular forces, efficiency of packing (molecular shape)
78
What are ionic solid properties
- Intermolecular Forces/Bonding: electrostatic attraction
between ions
- Physical Properties: hard, brittle, High melting and boiling points – dependant on surface charge density of ions
- Electrical and thermal insulators
- Some are soluble in polar solvents
79
What are metallic solid properties
- Intermolecular Forces/Bonding: Metal ions in a “sea” of valence electrons
- Physical Properties: Soft, ductile, malleable, Variable mp & bp, Many are soluble in (react with) acids, Electrical and thermal conductors
80
What are network covalent solids
- Intermolecular forces – Covalent bonding and van der waals force
- Physical Properties:Very hard & brittle, Very high bp & mp, Insulators or semiconductors, Insoluble in solvents
81
What are the 2 parts to a crystal
1. Lattice:
- Repeating 3-D pattern of points in space
- Theoretical
2. Basis
- Chemical unit (atom, molecule, etc.)
- One basis unit at each lattice point
82
what compounds typically form crystalline solids
Metals and ionic compounds
83
what compounds typically form amorphous solids
Substances that consist of large molecules, or a mixture of molecules whose movements are more restricted
84
What is a unit cell
The structure of a crystalline solid, whether a metal or not, is best described by considering its simplest repeating unit, which is referred to as its unit cell. The unit cell consists of lattice points that represent the locations of atoms or ions. The entire structure then consists of this unit cell repeating in three dimensions
