Liquids Chapter 10

Review: Gases Indefinite shape Indefinite volume Take the shape and volume of container Particles are far apart Particles move fast Low Density Easy to expand and compress

Review: Solids Definite shape Definite volume Particles close together, fixed Particles move very slowly High density Hard to expand/compress

Liquids: in between Closer to properties of solids Slow diffusion High attraction between particles Medium amount of energy

Forces of Attraction Intramolecular forces: Hold atoms together within a molecule covalent and ionic bonds Intermolecular forces: Hold molecules to each other 3 types

Dipole-Dipole Attraction Dipole: molecule with a separation of charge (polar covalent) Due to differences in electronegativity ~1% as strong as a covalent bond

Hydrogen Bond Very strong dipole-dipole attraction (10% of a covalent bond’s strength) Occurs when H is bonded to O, N, F in a very polar bond O H 3.5 2.1 Gives water its unusual properties

H-Bonding Affects Boiling Points Strong attraction requires much energy to overcome, so water is a liquid at normal temperatures

London Dispersion Forces Occur in all substances-polar and nonpolar Due to formation of instantaneous dipoles as electrons moving around nucleus concentrate on 1 side of molecule or atom

This induces a dipole in neighboring atoms or molecules These are the weakest intermolecular forces These are the only forces of attraction in nonpolar substances

Importance of Water Covers 70% earth’s surface Necessary for reactions in living cells Moderates earth’s temperature Coolant for engines & nuclear power plants Transportation Growth medium for many organisms

Properties of Water Colorless Tasteless At 1 atm, water freezes at 0°C and vaporizes completely at 100°C  Liquid phase occurs from 0-100°C

Special Properties of Water Surface Tension Liquids tend to form a “skin” making the surface less penetrable by solids

Unequal attraction at surface, mainly down Equal attraction in all directions

Surface Tension Detergents can interfere with the attractions and will cause the paperclip to fall

Adhesion Attraction of the surface of a liquid to the surface of a solid Depends on the material Water is attracted to glass Mercury is not No adhesion

Cohesion Molecular attractions within a material (water molecules to water molecules, for example) Here, cohesion causes water to form beads; ad- hesion causes it to stick to the web

Capillary Action A liquid rises in a narrow tube when it breaks the surface tension Movement of water through paper

Ice Floats! Molecules in a liquid have more movement than a solid and more energy (particles move apart Generally, solids are more dense than their ir liquids Liquid water

Ice When water becomes fixed points in a solid, hydrogen bonds hold molecules in place Gives ice an open hexagonal structure Greater volume means lower density than a liquid

Higher K.E. causes distance between molecules to be more Max. density at 4 C Lower K.E. causes distance between molecules to be less

Phase Changes of Water At 1 atm, water freezes at 0°C and vaporizes completely at 100°C  Liquid phase occurs from 0-100°C Changes from one phase to another will either require energy or release energy Solid Liquid Liquid Gas Solid Gas Melting/Freezing Vaporization/Condensation Sublimation/Deposition

From Solid to Liquid As energy is added, K.E. increases Solid warms up At 0°C, solid begins to melt and temperature remains at 0 until all solid is turned to liquid When all is liquid, temperature begins to rise

From Liquid to Gas As heat is added, K.E. increases (increase in temperature) At 100C, bubbles form in liquid Temperature remains the same until all liquid is converted to a gas. Once all is a gas, added energy causes temperature to increase

Ice melting/water vaporizing When a substance is in phase, increasing the energy increases the temperature When a substance is changing phase, increasing the energy does not increase the temperature but is used to break forces of attraction between molecules

Phase diagram Temperature vs. Energy

Heating a Solid

Melting a Solid

Heating a Liquid

Vaporizing a Liquid

Heating a Gas

Calculating Energies Energy is measured in calories or Joules 1.00 cal = 4.184 J The amount of energy needed to change states depends on: Type of matter Quantity of matter

Type of matter Molar heat of fusion (Hfusion)= energy needed to melt 1 mole of a substance Molar heat of vaporization (Hvap)= energy needed to vaporize 1 mole of a substance

For Water (Hfusion) = 80.0 cal/g = .334 kJ/g (H vap) = 540. cal/g = 2.26 kJ/g

Finding Energy in a Phase Change Change from a solid liquid q = mHfusion Change from a liquid gas q = mHvap q = energy (cal or J) m = mass (g) Hfusion =heat of fusion (cal/g) q = energy (cal or J) m = mass (g) Hvap =heat of vaporization

Example How much heat in calories is needed to melt 15.0 g of water? q = mHfusion 15.0 g water x 80.0 cal = 1.20 x 103 cal 1 g water

Energy and Being In Phase When all of a substance is in one phase, (e.g. all liquid), the amount of energy required to cause a temperature change depends on: type of substance amount of substance range of temperature change

Type of Matter Specific Heat (c): Energy required to change the temperature of 1 gram of a substance by 1 Celsius degree cg = specific heat of a gas cl = specific heat of a liquid cs = specific heat of a solid

For Water cg = .480 cal/gC or 2.01 J/gC cl = 1.00 cal/gC or 4.18 J/gC cs = .500 cal/gC or 2.09 J/gC

Finding Energy When In Phase q = mcT T = Tfinal – Tinitial q = energy (cal or J) m = mass (g) c = specific heat (cal/gC) T =change in temperature (C)

Example How much energy is required to heat 50.0 g of water from 20.0 C to 85.0 C? q = mcl T T = 85.0 – 20.0 = 65.0 C q = (50.0g)(1.00 cal/gC)(65.0 C) q = 3,250 cal

Phase changes

In Phase

Phase diagram Temperature vs. Energy

Phase Change Problems Draw graph. Mark start and stop points. Every corner means a new equation is needed. Flat sections will use q = mH ( no T means no slope). Find each energy (q1, q2, q3..). Add all energies to get the total energy.

Phase Changes Vaporization (evaporation): molecules of a liquid escape the liquid’s surface Requires energy to overcome intermolecular forces Maxwell Boltzman distribution Molecules with enough energy to evaporate

To evaporate, a particle must: Be at the surface Have sufficient energy Be moving in the right direction I’m free!! Moving in the wrong direction Not enough energy Not at surface

Evaporation produces Vapor Pressure A closed container with a vacuum has a liquid added to it. Molecules begin to evaporate Some particles are recaptured by the liquid Eventually rate of particles leaving = rate of being recaptured

Equilibrium Vapor Pressure When rate of evaporation = rate of condensation the pressure becomes constant (Equilibrium vapor pressure)

Vapor pressure and temperature As temperature increases, more particles evaporate

Vapor Pressure The vapor pressure of substances varies greatly, depending on the strength of the forces of attraction between particles.

Volatile liquids evaporate easily Small particle size Only London dispersion forces of attraction

Boiling Occurs when the equilibrium vapor pressure reaches atmospheric pressure Only then can a bubble maintain itself anywhere in the liquid Air Pressure Particle of gas exerts pressure on surrounding molecules, pushing them out of the way, and therefore up against air pressure

This is the difference between evaporation and boiling

As atmospheric pressure changes, so does the equilibrium vapor pressure necessary for boiling to occur Therefore, the boiling point depends on the pressure

Vapor Pressure vs. Temperature can give us the boiling point of a substance at any pressure The b.p. at 1 atm is called the normal boiling point.

P vs. T Phase Diagram We can expand our graph to see the effect of pressure on melting/freezing and sublimation/deposition For a “normal” substance: Increasing pressure raises the melting point (AD)

Important Points on Diagram A: Triple point: temperature & pressure at which solid, liquid and gas coexist at equilibrium B: Critical Point: indicates critical temp and pressure Critical Temp: temp above which substance cannot exist in a liquid state Critical Pressure: lowest pressure required for substance to exist as a liquid at the critical temperature

Phase Diagram for Water Line AB slants backwards. Increasing the pressure lowers the melting point.

1 atm 100°C 0°C