Unit 9: Liquids

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Unit 9 Liquids Solids
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Three States of Matter
State Shape Volume Why?
Gas None None Particles far apart
forces small
Liquid None Fixed Particles closer forces
greater
Solid Fixed Fixed Particles touch forces
great
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Why Gas Laws and Not Solid or Liquid Laws?

• Gases are mainly empty space have weak
  attractions between molecules
• Solids/liquids have particles which are closer
  together and have more varied forces between
  particles.

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Phase Transitions

• Definition
• Physical changes that result in changes of
  state.
• ALL phase changes involve energy (enthalpy).
• Review
• Enthalpy ?H heat at const P

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Phase Changes that Require Energy

• If you have to put energy into a reaction to make
  it happen, it is an endothermic reaction.
• Endothermic Phase Changes
• Melting (solid ? liquid) a.k.a fusion
• Vaporization (liquid ? gas)
• Sublimation (solid ? gas)

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Phase Changes that Release Energy

• If energy is released or given off by a reaction,
  it is an exothermic reaction.
• Exothermic Phase Changes
• Condensation (gas ? liquid)
• Deposition (gas ? solid)
• Freezing (liquid ? solid)

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NOTE What is constant at every phase change
above????
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Vaporization

• Definition conversion of liquid to gas.
• (endothermic)
• We commonly call this evaporation.
• Condensation is the reverse process
• (exothermic)

Evaporation
condensation
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Enthalpy of Vaporization

• Heat that is absorbed to vaporize a given amount
  of liquid at a constant temperature and pressure.
• Units Energy (J, kJ or cal) per g or mol
• LIQUID heat ---gt VAPOR
• Compd. ?Hvap (kJ/mol)
• H2O 40.7 (at 100 oC)
• SO2 26.8 (-47 oC)
• Xe 12.6 (-107 oC)

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Heat of Condensation is same quantity but
opposite in sign!

• ?Hvap - ?Hcond
• VAPOR ---gt LIQUID heat
• ?Hcond for water - 40.7 kJ/mol

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• Example 9. 1 How many kilojoules of heat are
  required to vaporize 598.5 g of ethanol? The heat
  of vaporization is 43.3 kJ/mol.

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Vapor Pressure

• Definition Pressure exerted by a vapor in a
  closed flask in equilibrium with its liquid
• Equilibrium two opposing processes occur at same
  rate.

A Rate vapgt Rate cond B Rate vap Rate cond
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• Example 9.2 Liquid ethanol has a vapor pressure
  of 43.9 mm Hg at 20 deg. C.
• What is the minimum volume of a flask needed to
  vaporize 1.00 g of liquid ethanol?

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Graphing Vapor Pressure
As Temperature increases, vapor pressure
increases. As attractive forces between molecules
increase, vapor pressure decreases. Attractive
Forces Water gt Ethanol gt Ether
Liquids with higher vapor pressure at a given T
are said to be more volatile.
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Questions1. What does this graph tell you
about the relative attraction between molecules
for substances a e? 2. Which substance is
most volatile?
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Boiling Point

• A liquid boils when
• vapor pressure atmospheric pressure.
• Normal Boiling Point
• Temp. where P 760 mm Hg (1 atm) on vapor
  pressure curve.
• Dependency on pressure
• As pressure increases, boiling point increases
  as pressure decreases, boiling point decreases.

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Normal Boiling Point
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Melting and Freezing Point

• Conversion of a solid to a liquid is called
  melting, or fusion.
• Conversion of a liquid to a solid is called
  freezing.
• The freezing point melting point
• Energy needed to melt a given quantity of solid
  is called the enthalpy of fusion, or ? H fus.

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• Example 11.3 How much energy is required to melt
  100.0 g of ice? The heat of fusion of water is
  6.01 kJ/mol.

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• Example 11.4 How much energy in kJ is required to
  heat 100.0 g of liquid water from zero to 100
  deg. C, and then vaporize all of it?
• Strategy
• 1. Calculate energy needed to heat water (Q
  equation from Chapter 6)
• 2. Calculate how much energy is needed to
  vaporize the water.
• 3. Add the two amounts together.

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• Example 11.4 How much energy in kJ is required to
  heat 100.0 g of liquid water from zero to 100
  deg. C, and then vaporize all of it?

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• Example 11.5 How much energy is required to heat
  75.0 g of ice from 0.0 deg. C to 185.0 deg. C?
  The heat capacity of steam is 1.84 J/g deg. C.
• Strategy
• 1. Calculate energy needed to melt ice.
• 2. Calculate how much energy is needed to heat
  the water to boiling.
• 3. Calculate how much energy is needed to
  vaporize the water.
• 4. Calculate how much energy is needed to
  heat the steam.
• 5. Add the four amounts together.

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• Example 11.5 How much energy is required to heat
  75.0 g of ice from 0.0 deg. C to 185.0 deg. C?
  The heat capacity of steam is 1.84 J/g deg. C.

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Cooling Curve H2O (g) ? H2O (s)
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Intramolecular Forces

• The attractive forces that hold particles
  together in ionic, covalent and metallic bonds
  are called intramolecular forces
• Intra- prefix within
• The forces inside a molecule holding the
  individual atoms together
• Ex.) Covalent bonds
• Ionic bonds
• Metallic bonds

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Intermolecular Forces (IMFs) Bonds between
Molecules

• Much weaker than Chemical Bonding within
  molecules
• Chemical Bonds (ionic and covalent) determine
  chemical properties
• Intermolecular forces determine physical
  properties
• e.g. density, mp, bp, solubility, vapor
  pressure, etc.

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Three General Types of Intermolecular Forces
(IMFs)

• Dispersion (London, van der Waals)
• Dipole/dipole
• Hydrogen Bonding

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London Dispersion Forces

• Definition IMF between two non-polar molecules
  formed by temporary positive and negative
  attractions due to the shifting of electron
  cloud.
• Found in all substances, but become important
  when they are the only IMF present.
• Strength increases as molar mass increases.

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London Dispersion Forces
Formation of a dipole in two nonpolar I2
molecules
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London Dispersion Forces

• Higher molar mass ---gt larger dispersion forces
• Molecule Boiling Point (oC)
• CH4 (methane) - 161.5
• C2H6 (ethane) - 88.6
• C3H8 (propane) - 42.1
• C4H10 (butane) - 0.5

Higher boiling point means GREATER IMFs!
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Example 11.6. Account for the fact that chlorine
is a gas, bromine is a volatile liquid, and
iodine is a volatile solid at room temperature.
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Dipole-Dipole Forces
Definition Attractions between oppositely
charged regions of polar molecules. Caused by
attraction of one dipole for another. Present
in all polar substances! Solubility and
dipole-dipole forces like dissolves like
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Dipole-Dipole Forces
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Hydrogen Bonding

• A special type of dipole-dipole force occurring
  only between molecules with a H atom bonded to
  either a F, O, or N atom.
• How to recognize F, O, or N directly bonded to H
• Two reasons why hydrogen bonds are stronger than
  dipole-dipole forces
• a. F, O, N very electronegative
• b. H is a small atom
• Hydrogen bonding is FON!

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Hydrogen Bonding
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Hydrogen Bonding in H2O

• H-bonding is especially strong in water because
• The OH bond is very polar
• There are 2 lone pairs on the O atom
• Accounts for many of waters unique properties.

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Hydrogen Bonding in H2O

• H bonds ---gt abnormally high specific heat
  capacity of water (4.184 J/gK)
• This is the reason water is used to put out
  fires, it is the reason lakes/oceans control
  climate, and is the reason thunderstorms release
  huge energy.

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Hydrogen Bonding

• H bonds
• lead to
• abnormally
• high
• boiling
• point of
• water.

See Screen 13.7
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Hydrogen Bonding in DNA

• Hydrogen bonding plays a key role in maintaining
  the double helix structure of DNA

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Example

• Identify the types of intermolecular forces
  present in compounds of
• Hydrogen Fluoride
• Pentane (C5H12)
• Hydrochloric Acid
• Ethanol (Ethyl Alcohol)

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Example

• Rank these substances in terms of increasing
  boiling point.
• N2 CCl4 CH3Cl NH3

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Liquids

• Viscosity a measure of the resistance of a
  liquid to flow
• The particles in a liquid are close enough
  together that their attractive forces slow their
  movement as they flow past one another
• The stronger the attractive forces
  (intermolecular forces), the more viscous the
  liquid is.
• As temperature increases, viscosity decreases.

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Liquids

• Surface tension an inward force that tends to
  minimize the surface area of a liquid
• A measure of the inward pull by particles in the
  interior
• The stronger the intermolecular forces, the
  higher the surface tension

In water, this is due mainly to hydrogen bonding!
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Liquids

• Surfactant any substance that interferes with
  the hydrogen bonding between water molecules
  reduces surface tension

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• Surfactants used to clean up oil spills as well
• Exxon Valdez oil spill in 1989 spilled over
  700,000 barrels of oil into the water near Alaska

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Network Covalent Solids

• Giant molecules connected by strong covalent
  bonds
• Properties hard, high mp, nonconductors

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Network Covalent Solids

• Atoms that can form multiple covalent bonds
• (look for C, Si, and other Group 14 elements)
• are able to form network covalent solids.
• All atoms in the entire structure are bonded
  together with covalent chemical bonds.

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Physical Properties of Graphite vs. Diamond
Property Graphite Diamond
Density (g/mL) 2.27 3.51
Hardness Very soft Very hard
Color Shiny black Colorless/transparent
Electrical Conductivity High None
DHcomb (kJ/mol) -393.5 -395.4
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Metallic and Ionic Solids See Chem Act. 24 and
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Ionic Solids

• A compound where each cation is simultaneously
  attracted to an anion.
• Review How can we identify an ionic compound?

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Properties of Ionic Solids

• 1. Molecules, atoms or ions locked into a
  CRYSTAL LATTICE
• 2. Particles are CLOSE together
• 3. STRONG IM forces
• 4. Highly ordered, rigid, incompressible

ZnS, zinc sulfide
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Metallic Solids

• A solid consisting of entirely metals.
• Characteristics
• Electrons are delocalized (they can move
  freely)
• Good conductors, malleable, ductile

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Metallic Solids

• Metallic solids positive metal ions surrounded
  by a sea of mobile electrons
• Mobile electrons make metals malleable and
  ductile because electrons can shift while still
  keeping the metal ions bonded in their new places
• Metallic solids are good conductors of heat and
  electricity
• Metallic Bonds

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Amorphous Solids

• Amorphous solid a solid in which the particles
  are not arranged in a regular, repeating pattern
• Amorphous without shape
• Often form when a molten material cools too
  quickly to allow enough time for crystals to form
• Common examples glass, rubber, many plastics

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Types of Solids

• TYPE EXAMPLE FORCE
• Ionic NaCl, CaF2, ZnS Ion-ion
• Metallic Na, Fe Metallic
• Molecular Ice, I2 Dipole Ind. dipole
• Network Diamond Extended Graphite covalent