Liquids

1
Liquids
2
Kinetic Theory

• Explains states of matter in terms of movement of
  particles.
• For liquids
• Molecular forces stronger than gases
  ?molecules closer than gases.
• Molecular forces weaker than solids
  ?molecules farther apart than solids.

3
Phase Changes

• Solid ? ? LOW Ek
• ?
• ?
• Temperature (Ek) increases
• The movement of particles ?
• Collisions between particles ?
• Distance between particles ?, due to increased
  force of collisions.
• ?
• ?
• Particles move apart sufficiently to enter
  another phase of matter -
• 
  Liquid.
• For pure substances the melting/freezing phase
  change takes place at a definite temperature.

4
Vapor Equilibrium
5
Vapor Equilibrium

• Represents a CLOSED SYSTEM
• Some particles lose Ek and change from a gas to a
  liquid.
• CONDENSATION
• Vapor Ek ? Liquid
• Some particles have/gain sufficient Ek to escape
  from liquid phase to become gaseous.
• EVAPORATION
• Liquid Ek ? Vapor (Gas)

6
Rate evaporation Rate condensation

• the system is in a state of dynamic equilibrium.
• The number of gas particles changing to liquid
  equals the number of liquid particles changing to
  gas.
• The vapor region of the system is saturated
  (contains as many particles possible at a
  specific temperature).

7
Equilibrium Vapor Pressure

• The pressure exerted by the vapor portion of the
  system on the liquid portion.
• Depends on
• Type of liquid
• Temperature
• Generally,
• the LOWER the vapor pressure the STRONGER the
  intermolecular forces of the particles on the
  system.

8
LeChateliers Principle

• If stress a change in the conditions of the
  system is applied to a system, the equilibrium
  moves in the direction that will reduce the
  stress.

9

• Pressure increases --- System tries to relieve
  the stress by reducing pressure.
• In the system initially Vapor ? Liquid Ek
• Pressure ? ? ? more liquid produced
• Volume ? ?? particles allowed to
  expand ?? more vapor prod.
  
• Temperature ? ?? Pressure ? ? ? more
  vapor produced

10
Solid to Liquid

• Solids become liquids at the melting point.
• Temperature at which vapor pressure of the solid
  vapor pressure of liquid
• Depends on
• Size of the molecules
• Molecular forces
• Weak intermolecular forces --- LOW m.p.
• High intermolecular forces --- HIGH m.p.
• Ex H2O CH4
• Polar Nonpolar
• m.w. 18 m.w. 16
• m.p. 0o C m.p. 183 oC
• b.p. 100 0C b.p. 162 oC

11

• Standard Boiling Point
• Temperature _at_ which the vapor pressure of the
  liquid equals standard atmospheric pressure.
• 1 ATM, 101.325 kPa, 760 mm Hg, 760 Torr
• Boiling Point
• Temperature _at_ which the vapor pressure of the
  liquid equals the existing atmospheric pressure.

12

• If a liquid
• Boils quickly
• Evaporates quickly _at_ room temperature
• the liquid is said to be volatile.
• Volatile liquids usually have HIGH vapor
  pressures.
• Ex alcohols
• If the vapor pressure is extremely high the
  substance will change from a solid directly to a
  gas.
• Sublimation

13
Pascals Law

• Applies the idea of vapor pressure to BOILING
• Pressure is exerted evenly on the surface of a
  confined liquid, and is transferred undiminished
  throughout.

14
As a liquid is heated

• Gases become less soluble _at_ higher temperatures.
• ? bubbles form at the hottest point
• Bubbles collapse due to atmospheric pressure
  Pascals Law
• As the temperature of the liquid increases, the
  Ek of the molecules increases.
• ?the vapor pressure of the bubbles increases
• When the vapor pressure of the bubbles equals the
  
• vapor pressure of the atmosphere the bubbles
  survive to the surface of the liquid BOILING
• During Boiling the temperature of the liquid
  remains constant.

15
Why can steam burn more severely than boiling
water?

• A higher Ek

16
Liquification of Gases

• Michael Faraday
• Discovered it was possible to liquefy gases by
  simultaneously cooling and compressing the gas.

17
Modern Liquification of Gases

• Compress gas and increase temperature
• Use a coolant to remove the temperature increase
• Allow the gas to expand ? Temperature drops
  Joule-Thomas Effect
• Process repeated
• As temperature decreases ? Ek decreases
• A point is reached where the intermolecular
  forces of attraction van der Waal / London
  forces can cause the molecules to combine if
  they are close enough.

18
For every gas

• there is a temperature above which NO AMOUNT of
• pressure will cause the gas to liquefy critical
  temperature Tc.
• Pressure that results in liquification at
  Tc critical pressure.
• In order to liquefy a gas
• Conditions must be at / below Tc
• Pressure at / above vapor pressure of liquid

19
Energy and Changes of State
Ep inc.
Ek inc
Ep inc.
Ek inc
Before a phase change Kinetic Energy
increases. During a phase change Potential
Energy increases.
20
Enthalpy

• Enthalpy of Fusion Hfus heat energy required
  to melt 1 gram
• 
  of a substance.
• Ex ice 334 J/g
• Enthalpy of Vaporization Hvap heat energy
  required to convert
• 1 gram of liquid to gas _at_ its
  boiling
• 
  point.
• Ex H2O 2260 J/g
• Enthalpy q heat transferred
• q m?Hfus
• q m?Hvap
• q m?TCp specific heat

21
WATER

• Most abundant liquid
• 75 of Earths surface covered by it.
• 10 of the land covered by water in the form of
  glaciers
• Water vapor is always present in the air.
• 70 -90 of all living things are water by
  weight.
• Large reserves of water underground

22
Physical Properties

• Transparent
• Odorless
• Tasteless
• Almost colorless
• Exists as solid, liquid, gas
• _at_ Standard Pressure 1 Atm., 760 mm Hg, 760
  Torr, 101.325 kPa
• 0o C ----- liquid water ---------? solid water
  Energy
• 100 cm3 111.111
  cm3
• 1/9th greater volume
• 0o C ? 4o C --------- Water contracts
• gt 4o C ---------------- Water begins to expand
  
• Water most dense _at_ 4o C.
• 1 g 1 ml 1cm3

23
Standard Values for Water

• Melting / freezing ..0o C
• Boiling / condensing 100o C
• Molar heat of vaporization ...9.7 kcal
• Heat required to convert 1 mole of water (18 g)
  into a gas

24

• From the definition of boiling, we can see that
  as the pressure ? the boiling point ?.
• Putting this to practical application..
• Vacuum Evaporator
• Industrial device that reduces the b.p. by
  reducing the
• pressure on the liquid.
• Evaporated milk
• Eagle Brand milk
• Soups

25
Structure and Properties of Water

• Bent Molecule
• -
• O
• H H
• 104.5o

• Oxygen and hydrogen form a polar covalent bond
  dipole

26

• In a container
• Equal forces in all directions
• On the surface
• Perpendicular force downward on the surface
• The perpendicular force causes the phenomena of
• Surface tension
• Pulls surface molecules together and make it less
  
• penetrable than expected.
• Capillary rise
• Rise of a liquid in a tube of small diameter.
• The attraction between the water and the glass
  relieves the perpendicular force.

27

• For many substances the
• 1) structure and the
• 2) molecular weight can be used to predict the
  general changes in state.

28

• Generally _at_ room temperature
• Solids ionic bonds m.p. NaCl 801o C
• 58.5
• C6BrCH3 288o C
• Toluene
• 487
• Liquids Polar covalent bonds C2H4 -169o C
• Gases Pure covalent bonds N2 -210o C

29
Some substances do not change their state within
the ranges expected

• Contain H in a polar covalent bond
• H bonded to atoms like N, O, or F
• These atoms are highly electronegative

30

• Electronegative atoms pull shared electrons
  closer to them than to the other atom in the bond.

-
Hydrogen
Nitrogen
Shared e- Pair
The uneven sharing of the electron pair causes
slight charges (dipoles) to form around the
bond.
31
The Hydrogen Bond

• A weak attractive force not a bond between the
  hydrogen of one molecule and an electronegative
  atom in another molecule.

Hydrogen
Electronegative atom
Hydrogen bond
32

• Hydrogen bonds are sensitive to temperature.
  
• The higher the temperature the fewer the hydrogen
  bonds.
• The effect of hydrogen bonding is cumulative.
  Like pieces of paper/ threads/hair
• Hydrogen bonding causes water to be a liquid _at_
  room temperature.
• It would otherwise be a gas due to its low
  molecular weight.

33

• 0o C O H bonds between the water
• 
  molecules are rigid.
• As the temperature increases.
• The O H bonds become flexible.
• This flexibility allows the molecules to move
  closer
• together.
• 4o C Maximum density.
• At Higher Temperature Molecules begin to move
  apart.
• Liquid expands
• Hydrogen bonding explains why water expands when
  it freezes.

34
Chemical Behavior of Water

• Extremely stable molecule
• Does not begin to decompose until temperature is
  gt 2700o C.
• Reacts with metals to form hydrogen gas.
• H2O X ? H2 XOH
• Hydrogen gas Metal hydroxide _at_ rm
  temp
• H2O X ? H2 XOH
• Hydrogen gas Metal oxide _at_ temp
  gt 170oC
• Exception K can react with cold water to form a
  metal oxide.

35
Basic anhydride

• a metal oxide that forms a basic solution when
  combined with water
• BASIC solution due to OH-
• Slippery
• Bitter taste
• Red litmus ? Blue
• Reacts with metal oxides to form metal
  hydroxides.
• H2O XO? XOH
• XOH ? X OH-
• The hydroxide separates - forms the metal in
  solution X and OH- result.

36
Acid Anhydride

• a nonmetal oxide that forms an acid solution when
  combined with water
• ACID solution due to H3O
• Sour taste
• Blue litmus ? Red
• Reacts with nonmetal oxides to form acidic
  solutions.
• H2O CO2 ? H2CO3 carbonic acid
• The hydrogen separates form the nonmetal in
  solution H and HCO3result.
• H reacts with the water to form the hydronium
  ion. H3O

37
Capturing and Releasing Water

• Na2CO3 ? 10H2O ? Na2CO3 ? H2O 9H2O
• Loosely
• attached
• Water of hydration

38

• You go to the grocery store and buy a box of
  washing soda that is marked 14 oz. You weigh the
  box when you get home. The balance reads 10 oz.
• Assuming that the company is honest and that no
  mistakes were made in packaging, how can you
  explain this discrepancy?

39

• Efflorescence
• Water of hydration that is lost to the
  atmosphere.
• The higher the vapor pressure of the substance
  the greater the efflorescence of the hydrate.
• Deliquescence
• Absorption of water molecules from the
  atmosphere.
• The vapor pressure of the substance is lower than
  the vapor pressure of the water vapor in the air.
• Forms a solution with the water in the air until
  the vapor pressure of each is equal.
• Hydroscopic
• Takes water form atmosphere and traps it in pores
  and imperfections in its surface.
• Ex hair, wool, tobacco, potato chips