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AP Chapter 11. Intermolecular Forces / Liquids and Solids HW: 2 4 6 13 15 16 19 31 39 45 105. 11.1 – A Comparison of the States of Matter. The fundamental difference between states of matter is the distance between particles. Indefinite volume Indefinite shape

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ap chapter 11

AP Chapter 11

Intermolecular Forces / Liquids and Solids

HW: 2 4 6 13 15 16 19 31 39 45 105

11 1 a comparison of the states of matter
11.1 – A Comparison of the States of Matter

The fundamental difference between states of matter is the distance between particles.

Indefinite volume

Indefinite shape

Rotational, Vibrational and Translational Particle Motion

Definite volume

Indefinite shape

Rotational and Vibrational

Definite volume

Definite shape

Rotational Motion

states of matter
States of Matter

Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.

phase
Phase
  • Homogeneous part of a system in contact with other parts of the system but separated from them by a well defined boundary
  • The state/phase a substance is in at a particular temperature and pressure depends on two properties:
    • The kinetic energy of the particles
    • The strength of the attractions between the particles
11 2 intermolecular forces external bonds
11.2 - Intermolecular Forces (External Bonds)

The attractions between molecules are not nearly as strong as the intramolecular attractions (covalent bonds) that hold compounds together.

intermolecular forces
Intermolecular Forces

They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

1 ion dipole interactions
1. Ion-Dipole Interactions
  • Ion-dipole interactions are an important force in solutions of ions.
  • The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents
  • Hydration = Process of water molecules surrounding a solute
2 intermolecular forces
2. Intermolecular Forces

Intermolecular forces between molecules are referred to as van der Waals forces.

van der waals forces
van der Waals Forces
  • a. Dipole-dipole interactions
  • b. London Dispersion forces
    • Induced Dipoles (not a true London Dispersion Force)
    • Instantaneous Dipoles
a dipole dipole interactions
a. Dipole-Dipole Interactions
  • Molecules that have permanent dipoles are attracted to each other.
    • The positive end of one is attracted to the negative end of the other and vice-versa.
    • These forces are only important when the molecules are close to each other.
    • The larger the dipole moment, the stronger the attraction
dipole dipole interactions
Dipole-Dipole Interactions

The more polar (larger resultant dipole moment) the molecule, the higher is its boiling point because the attractions are stronger.

hydrogen bonding
Hydrogen Bonding
  • Can be considered a unique kind of dipole-dipole interaction
  • Hydrogen atom in a polar bond externally bonds to the non-bonding electron pair (lone pair) on an electronegative atom
    • Polar bond usually the H connected to a N, O, or F
  • We call these interactions hydrogen bonds.
  • Often seen in biological systems.
hydrogen bonding1
Hydrogen Bonding

Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.

Because when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

slide14

b. London Dispersion Forces(occur between nonpolar substances/atoms)-Induced Dipoles (not all books talk about these. Not a TRUE London Dispersion Force)

  • Occurs when a nonpolar atom/molecule comes in proximity to charged particle/molecule
  • Induces a charge on the nonpolar molecule

And then is attracted to dipole

+

->

london dispersion force instantaneous dipole
London Dispersion Force-Instantaneous Dipole

While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom. Causes an instantaneous dipole. The same can occur in any molecule…

instantaneous dipole
Instantaneous Dipole

Looks like this

At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

london dispersion forces
London Dispersion Forces

Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

london dispersion forces1
London Dispersion Forces

-A London Dispersion Force Occurs

-London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

london dispersion forces2
London Dispersion Forces
  • These forces are present in all molecules, whether they are polar or nonpolar.
  • The tendency of an electron cloud to distort in this way is called polarizability.
factors affecting london forces
Factors Affecting London Forces
  • The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).
  • This is due to the increased surface area in n-pentane.

(that’s why oil is thicker than water)

factors affecting london forces1
Factors Affecting London Forces
  • The strength of dispersion forces tends to increase with increased molecular weight.
  • Larger atoms have larger electron clouds, which are easier to polarize.
dipole dipole interactions vs london forces
Dipole-Dipole Interactions vs. London Forces
  • Nonpolar molecules only have London Forces
  • Dipole molecules have both Dipole-Dipole and London Forces
    • Which is more important to the property of the molecule??
  • When comparing two dipole molecules with different BP, etc.:
    • If size, shape and # of e-s of the two molecules are very similar, the reason the BPs are different are due to dipole-dipole interactions
    • If the molecules are of very different size than dispersion forces are the main explanation
slide23

Dispersion Forces vs. Dipole-Dipole Interactions

Why choose these two molecules to compare? Both have identical numbers of electrons, and if you made models you would find that the sizes were similar - as you can see in the diagrams. That means that the dispersion forces in both molecules should be much the same.

The higher boiling point of fluoromethane is due to the large permanent dipole on the molecule because of the high electronegativity of fluorine. However, even given the large permanent polarity of the molecule, the boiling point has only been increased by some 10°. (Dipole-Dipole Interactions don’t always make that much of a difference).

dispersion forces vs dipole dipole interactions

The boiling points are:

Dispersion Forces vs. Dipole-Dipole Interactions

Here is another example showing the dominance of the dispersion forces. Trichloromethane, CHCl3, is a highly polar molecule because of the electronegativity of the three chlorines. There will be quite strong dipole-dipole attractions between one molecule and its neighbours.

On the other hand, tetrachloromethane, CCl4, is non-polar. The outside of the molecule is uniformly - in all directions. CCl4 has to rely only on dispersion forces

So which has the highest boiling point? CCl4 does, because it is a bigger molecule with more electrons. The increase in the dispersion forces more than compensates for the loss of dipole-dipole interactions.

11 3 physical properties of liquids
11.3 - Physical Properties of Liquids

The strength of the attractions between particles can greatly affect the properties of a substance or solution.

viscosity
Viscosity
  • Resistance of a liquid to flow is called viscosity.
  • It is related to the ease with which molecules can move past each other.
  • Viscosity increases with stronger intermolecular forces and decreases with higher temperature.
surface tension
Surface Tension

No molecules pull upward on the surface

Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

capillary action
Capillary Action
  • Liquid in a tube is not at the same level as the liquid in the surrounding container

-Cohesion – Intermolecular attraction between liquid molecules

-Adhesion – Attraction between unlike molecules (ex – liquid and the tube)

Adhesion > Cohesion

Cohesion > Adhesion

11 4 phase changes
11.4 – Phase Changes
  • Transfer from one phase to another by addition or removal of energy
  • Liquid – Vapor Equilibrium
    • When molecules have enough energy they escape the surface
    • Higher T, Higher E, More Evaporation
vapor pressure
Vapor Pressure
  • At any temperature, some molecules in a liquid have enough energy to escape.
  • As the temperature rises, the fraction of molecules that have enough energy to escape increases.
vapor pressure1
Vapor Pressure

As more molecules escape the liquid, the pressure they exert increases.

vapor pressure2
Vapor Pressure

The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate. (in a closed container)

vapor pressure3
Vapor Pressure
  • The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure.
  • The normal boiling point is the temperature at which its vapor pressure is 760 torr.
  • Molar Heat of Vaporization = Energy needed to vaporize one mole of liquid.
clausius clapeyron equation p 456
Clausius-Clapeyron Equation (p. 456)
  • Describes the relationship between P (vapor pressure) and T
  • Plot ln P on y-axis vs 1/T on x-axis
    • Slope = -DHvap/R
    • y-intercept = c
    • R = 8.314 J/molK
  • Comparing two Temps:

ln(Pvap) = -DHvap/(RT) + c

Sometimes shown without – and T2 and T1 reversed…

boiling point
Boiling Point
  • Temp at which VP liquid = P external
  • Water = Normal BP = 100oC
    • Normal = P external = 1 atm
critical temperature and pressure
Critical Temperature and Pressure
  • Gas -> Liquid occurs by either decrease temp or increase pressure
  • Critical Temp = Tc = “Last” temperature at which a liquid can exist. At ALL pressures above this temp, only a gas can exist
  • Critical Pressure = Pc = Minimum pressure that must be applied to bring about liquification at the Critical Temperature
liquid solid equilibrium
Liquid – Solid Equilibrium

Heating / Cooling Curves:

  • Freezing = l->s
  • Melting/Fusion = s->l
  • Melting Point / Freezing Point = Temperature at which solid and liquid are at equilibrium
  • Water = Normal fp = 0oC
  • Molar Heat of Fusion = Energy required to melt 1 mole of a solid
supercooling
Supercooling
  • Liquid temporarily cooled below Freezing Point
  • Happens when removed from heat so rapidly that molecules don’t have time to form solid.
    • Unstable
    • Can solidify by seed crystal or scratch on side of container
solid vapor equilibrium
Solid-Vapor Equilibrium
  • Sublimation = s->g
  • Deposition = g->s
  • Molar Heat of Sublimation = Energy required to sublime 1 mole solid
heat calculations
Heat Calculations
  • BEFORE Phase Changes (temp is changing):

q = msDT

  • DURING Phase Changes (temp constant due to all energy going into changing bonds):

q = m*1/M*DH

moles

11 6 phase diagrams
11.6 - Phase Diagrams

Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

phase diagrams
Phase Diagrams
  • Triple Point = Conditions at which s, l, g all coexist
    • Water = 0.01oC and 0.006 atm
  • The AB line is the liquid-vapor interface.
  • It starts at the triple point (A), the point at which all three states are in equilibrium.
phase diagrams1
Phase Diagrams

It ends at the critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

phase diagrams2
Phase Diagrams

Each point along this line is the boiling point of the substance at that pressure.

phase diagrams3
Phase Diagrams
  • The AD line is the interface between liquid and solid.
  • The melting point at each pressure can be found along this line.
phase diagrams4
Phase Diagrams
  • Below A the substance cannot exist in the liquid state.
  • Along the AC line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.
phase diagram of water
Phase Diagram of Water
  • Note the high critical temperature and critical pressure:
    • These are due to the strong van der Waals (both dipole-dipole and London Dispersion) forces between water molecules.
structure and properties of water
Structure and Properties of Water
  • Water is a very unique compound
  • Has a VERY high specific heat (can absorb large amounts of heat but with small change in temperature)
  • Solid is less dense than the liquid
  • Maximum density occurs at 4oC
phase diagram of water1
Phase Diagram of Water
  • The slope of the solid–liquid line is negative.
    • This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.
    • RARE
phase diagram of carbon dioxide
Phase Diagram of Carbon Dioxide

Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.

phase diagram of carbon dioxide1
Phase Diagram of Carbon Dioxide
  • Solid – Liquid line has a + slope
  • Cannot liquify a solid using pressure

*ice skating

11 7 crystal structure of solids
11.7 – Crystal Structure of Solids
  • We can think of solids as falling into two groups:
    • Crystalline—particles are in highly ordered arrangement.
solids
Solids
  • Amorphous—no particular order in the arrangement of particles.
crystalline solids
Crystalline Solids

Because of the order in a crystal, we can focus on the repeating pattern of arrangement called the unit cell.

unit cells
Unit Cells
  • There are 7 basic types of unit cells
11 8 bonding in solids molecular crystals
11.8 – Bonding in SolidsMolecular Crystals
  • Molecules held together by external bonds
    • Ex: Ice, Iodine (s), S8, P4
11 8 bonding in solids covalent network
11.8 – Bonding in SolidsCovalent-Network
  • Diamonds are an example of a covalent-network solid in which MANY atoms are covalently bonded to each other.
covalent network
Covalent-Network
  • Graphite is an example in which atoms are held together by covalent bonds and the layers are held with van der Waals forces.
11 8 bonding in solids ionic
11.8 – Bonding in SolidsIonic
  • Ions at the lattice pts
  • Cations and anions in alternating pattern

CaF2

CsCl

ZnS

11 8 bonding in solids metallic
11.8 – Bonding in SolidsMetallic
  • Metal Ions in an electron sea
slide62
Amorphous Solids

-No pattern

-Glass

-Quickly cooled liquid->solid