Liquids

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.

Phase Changes

       

Solid

  LOW E k   Temperature (E k ) increases The

movement Collisions Distance

of particles  between particles 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.

Vapor Equilibrium

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)

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).

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.

LeChatelier’s 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.

 Pressure increases --- System tries to relieve the stress by reducing pressure.

    In the system initially: Vapor ↔ Liquid + E k

Pressure

↑   more liquid produced

Volume

↑

Temperature

↓   particles allowed to expand 

Pressure

↓   more vapor prod. more vapor produced

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:

H 2 O

Polar m.w. 18 m.p. 0o C b.p. 100 0C

CH 4

Nonpolar m.w. 16 m.p. –183 oC b.p. –162 oC



Standard Boiling Point

 Temperature @ which the vapor pressure of the liquid equals standard atmospheric pressure.

 [1 ATM, 101.325 kPa, 760 mm Hg, 760 Torr] 

Boiling Point

 Temperature @ which the vapor pressure of the liquid equals the existing atmospheric pressure.

    If a liquid    Boils quickly Evaporates quickly @ 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

Pascal’s 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.

As a liquid is heated…

       Gases become less soluble @ higher temperatures.

 bubbles form at the hottest point Bubbles collapse due to atmospheric pressure [Pascal’s 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.

Why can steam burn more severely than boiling water?

 A: higher E k

Liquification of Gases



Michael Faraday

 Discovered it was possible to liquefy gases by

simultaneously

cooling and compressing the gas.

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.

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

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.

Enthalpy



Enthalpy of Fusion

[H fus ] – heat energy required to melt 1 gram of a substance.

 Ex: ice – 334 J/g  

Enthalpy of Vaporization

Ex: H2O – 2260 J/g [H vap ] – heat energy required to convert 1 gram of liquid to gas @ its boiling point.



Enthalpy

   [q] – heat transferred q = m  H fus q = m  H vap q = m  TC p [specific heat]

WATER

 Most abundant liquid  75% of Earth’s 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

Physical Properties

        Transparent Odorless Tasteless Almost colorless Exists as solid, liquid, gas @ Standard Pressure [1 Atm., 760 mm Hg, 760 Torr, 101.325 kPa]   0 o C ----- liquid water --------  100 cm 3 solid water + Energy 111.111 cm 3 0 o C  4 o C -------- 1/9th greater volume Water contracts    > 4 o C ---------------- Water begins to expand Water

most dense @ 4 o C.

1 g = 1 ml = 1cm 3

Standard Values for Water

 Melting / freezing ……………..0

o C  Boiling / condensing ……………100 o C  Molar heat of vaporization ……...9.7 kcal  [Heat required to convert 1 mole of water (18 g) into a gas]

 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

Structure and Properties of Water

 Bent Molecule  O + H H+ 104.5

o •Oxygen and hydrogen form a

polar covalent

bond [dipole]

 

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.

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

 Generally @ room temperature…   

Liquids

Polar covalent bonds



Solids Gases

ionic bonds

m.p.

Pure covalent bonds

NaCl [58.5] 801 o C C 6 BrCH 3 288 o Toluene [487] C C N 2 2 H 4 -169 o C -210 o C

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

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

+ Hydrogen

Shared e- Pair Nitrogen The uneven sharing of the electron pair causes slight charges (

dipoles

) to form around the bond.

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”

 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 @ room temperature.  [It would otherwise be a gas due to its low molecular weight.]



0 o C O H bonds between

the water molecules are

rigid.

         As the temperature increases….

together.

4 o C

The This flexibility allows the

O H

bonds become

Maximum density.

flexible.

molecules to move closer

 At Higher Temperature Molecules begin to move apart.  Liquid expands Hydrogen bonding explains why water expands when it freezes.

Chemical Behavior of Water

       Extremely stable molecule Does not begin to decompose until temperature is > 2700 o C.

 Reacts with metals to form hydrogen gas.

H 2 O + X



H 2 + XOH

Hydrogen gas + Metal hydroxide @ rm temp

H 2 O + X



H 2 + XOH

Hydrogen gas + Metal oxide @ temp > 170 o C [Exception: K can react with cold water to form a metal oxide.]

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.

H 2 O + XO



XOH XOH



X + + OH -

The hydroxide separates - forms the metal in solution [ X+ and OH- result].

Acid Anhydride

      a nonmetal oxide that forms an acid solution when combined with water ACID solution [due to H 3 O + ]   Sour taste Blue litmus  Red Reacts with nonmetal oxides to form acidic solutions.

H 2 O + CO 2



H 2 CO 3

[carbonic acid] The hydrogen separates form the nonmetal in solution [ H + and HCO 3 —result].

H+ reacts with the water to form the

hydronium ion.

[

H 3 O + ]

Capturing and Releasing Water

 Na 2 CO 3  10H 2 O  Na 2 CO 3  H 2 O + 9H 2 O Loosely attached Water of hydration

 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?

 

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