Chapter 12 Liquids and Solids

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Chapter 12LiquidsandSolids
2006, Prentice Hall
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Interactions Between Molecules

• many of the phenomena we observe are related to
  interactions between molecules that do not
  involve a chemical reaction
• your taste and smell organs work because
  molecules in the thing you are sensing interact
  with the receptor molecule sites in your tongue
  and nose
• in this chapter we will examine the physical
  interactions between molecules and the factors
  that effect and influence them

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The Physical States of Matter

• matter can be classified as solid, liquid or gas
  based on what properties it exhibits

• Fixed keeps shape when placed in a container,
• Indefinite takes the shape of the container

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Structure Determines Properties

• the atoms or molecules have different structures
  in solids, liquid and gases, leading to different
  properties

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Properties of the States of MatterGases

• low densities compared to solids and liquids
• fluid
• the material exhibits a smooth, continuous flow
  as it moves
• take the shape of their container
• expand to fill their container
• can be compressed into a smaller volume

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Properties of the States of MatterLiquids

• high densities compared to gases
• fluid
• the material exhibits a smooth, continuous flow
  as it moves
• take the shape of their container
• keep their volume, do not expand to fill their
  container
• can not be compressed into a smaller volume

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Properties of the States of MatterSolids

• high densities compared to gases
• nonfluid
• they move as entire block rather than a smooth,
  continuous flow
• keep their own shape, do not take the shape of
  their container
• keep their own volume, do not expand to fill
  their container
• can not be compressed into a smaller volume

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The Structure of Solids, Liquid and Gases
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Gases

• in the gas state, the particles have complete
  freedom from each other
• the particles are constantly flying around,
  bumping into each other and the container
• in the gas state, there is a lot of empty space
  between the particles
• on average

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Gases

• because there is a lot of empty space, the
  particles can be squeezed closer together
  therefore gases are compressible
• because the particles are not held in close
  contact and are moving freely, gases expand to
  fill and take the shape of their container, and
  will flow

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Liquids

• the particles in a liquid are closely packed, but
  they have some ability to move around
• the close packing results in liquids being
  incompressible
• but the ability of the particles to move allows
  liquids to take the shape of their container and
  to flow however they dont have enough freedom
  to escape and expand to fill the container

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Solids

• the particles in a solid are packed close
  together and are fixed in position
• though they are vibrating
• the close packing of the particles results in
  solids being incompressible
• the inability of the particles to move around
  results in solids retaining their shape and
  volume when placed in a new container and
  prevents the particles from flowing

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Solids

• some solids have their particles arranged in an
  orderly geometric pattern we call these
  crystalline solids
• salt and diamonds
• other solids have particles that do not show a
  regular geometric pattern over a long range we
  call these amorphous solids
• plastic and glass

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Why is Sugar a Solid ButWater is a Liquid?

• the state a material exists in depends on the
  attraction between molecules and their ability to
  overcome the attraction
• the attractive forces between ions or molecules
  depends on their structure
• the attractions are electrostatic
• depend on shape, polarity, etc.
• the ability of the molecules to overcome the
  attraction depends on the amount of kinetic
  energy they possess

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Properties of LiquidsViscosity

• some liquids flow more easily than others
• the resistance of a liquid to flow we call
  viscosity
• larger the attractive forces between the
  molecules larger the viscosity
• also, molecules whose shape is not round will
  have a larger viscosity

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Properties of LiquidsSurface Tension

• liquids tend to minimize their surface a
  phenomenon we call surface tension
• this tendency causes liquids to have a surface
  that resists penetration

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Surface Tension

• molecules in the interior of a liquid experience
  attractions to surrounding molecules in all
  directions
• but molecules on the surface experience an
  imbalance in attractions, effectively pulling
  them in
• to minimize this imbalance and maximize
  attraction, liquids try to minimize the number of
  molecules on the exposed surface by minimizing
  their surface area
• stronger attractive forces between the molecules
  larger surface tension

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Forces of Attraction within a Liquid

• Cohesive Forces forces that try to hold the
  liquid molecules to each other
• surface tension
• Adhesive Forces forces that bind a substance to
  a surface
• capillary action
• meniscus

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Escaping from the Surface

• the process of molecules of a liquid breaking
  free from the surface is called evaporation
• also known as vaporization
• evaporation is a physical change in which a
  substance is converted from its liquid form to
  its gaseous form
• the gaseous form is called a vapor

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Evaporation

• over time, liquids evaporate the molecules of
  the liquid mix with and dissolve in the air
• the evaporation happens at the surface
• molecules on the surface experience a smaller net
  attractive force than molecules in the interior
• but all the surface molecules do not escape at
  once, only the ones with sufficient kinetic
  energy to overcome the attractions will escape

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Factors Effecting the Rate of Evaporation

• increasing the surface area increases the rate of
  evaporation
• increasing the temperature increases the rate of
  evaporation
• weaker attractive forces between the molecules
  faster rate of evaporation
• liquids that evaporate quickly are called
  volatile liquids, while those that do not are
  called nonvolatile

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Escaping the Surface

• the average kinetic energy is directly
  proportional to the kelvin temperature
• but not all molecules in the sample have the same
  kinetic energy
• those molecules on the surface that have enough
  kinetic energy will escape
• raising the temperature increases the number of
  molecules with sufficient energy to escape

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Escaping the Surface

• since the higher energy molecules from the liquid
  are leaving, the total kinetic energy of the
  liquid decreases, and the liquid cools
• the remaining molecules redistribute their
  energies, generating more high energy molecules
• the result is the liquid continues to evaporate

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Reconnecting with the Surface

• when a liquid evaporates in a closed container,
  the vapor molecules are trapped
• the vapor molecules may eventually bump into and
  stick to the surface of the container or get
  recaptured by the liquid this process is called
  condensation
• a physical change in which a gaseous form is
  converted to a liquid form

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Dynamic Equilibrium

• evaporation and condensation are opposite
  processes
• eventually, the rate of evaporation and
  condensation in the container will be the same
• opposite processes that occur at the same rate in
  the same system are said to be in dynamic
  equilibrium

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Evaporation and Condensation
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Vapor Pressure

• once equilibrium is reached, from that time
  forward, the amount of vapor in the container
  will remain the same
• as long as you dont change the conditions
• the partial pressure exerted by the vapor is
  called the vapor pressure
• the vapor pressure of a liquid depends on the
  temperature and strength of intermolecular
  attractions

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Boiling

• in an open container, as you heat a liquid the
  average kinetic energy of the molecules
  increases, giving more molecules enough energy to
  escape the surface
• so the rate of evaporation increases
• eventually the temperature is high enough for
  molecules in the interior of the liquid to escape
  a phenomenon we call boiling

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Boiling Point

• the temperature at which the vapor pressure of
  the liquid is the same as the atmospheric
  pressure is called the boiling point
• the normal boiling point is the temperature
  required for the vapor pressure of the liquid to
  be equal to 1 atm
• the boiling point depends on what the atmospheric
  pressure is
• the temperature of boiling water on the top of a
  mountain will be cooler than boiling water at sea
  level

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Temperature and Boiling

• as you heat a liquid, its temperature increases
  until it reaches the boiling point
• once the liquid starts to boil, the temperature
  remains the same until it all turns to a gas
• all the energy from the heat source is being used
  to overcome the attractive forces in the liquid

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Energetics of Evaporation

• as it loses the high energy molecules through
  evaporation, the liquid cools
• then the liquid absorbs heat from its
  surroundings to raise its temperature back to the
  same as the surroundings
• processes in which heat flows into a system from
  the surroundings are said to be endothermic
• as heat flows out of the surroundings, it causes
  the surroundings to cool
• as alcohol evaporates off your skin, it causes
  your skin to cool

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Energetics of Condensation

• as it gains the high energy molecules through
  condensation, the liquid warms
• then the liquid releases heat to its surroundings
  to reduce its temperature back to the same as the
  surroundings
• processes in which heat flows out of a system
  into the surroundings are said to be exothermic
• as heat flows into the surroundings, it causes
  the surroundings to warm

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Heat of Vaporization

• the amount of heat needed to vaporize one mole of
  a liquid is called the heat of vaporization
• DHvap
• it requires 40.7 kJ of heat to vaporize one mole
  of water at 100C
• endothermic
• DHvap depends on the initial temperature
• since condensation is the opposite process to
  evaporation, the same amount of energy is
  transferred but in the opposite direction
• DHcond -DHvap

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Heats of Vaporization of Liquidsat their Boiling
Points and at 25C
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ExampleCalculate the amount of water in grams
that can be vaporized at its boiling point with
155 kJ of heat.

• Information
• Given 155 kJ
• Find g H2O
• CF 40.7 kJ 1 mol 18.02 g 1 mol
• SM kJ ? mol ? g

• Apply the Solution Map

68.626 g H2O

• Sig. Figs. Round

68.6 g H2O
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Temperature and Melting

• as you heat a solid, its temperature increases
  until it reaches the melting point
• once the solid starts to melt, the temperature
  remains the same until it all turns to a liquid
• all the energy from the heat source is being used
  to overcome the attractive forces in the solid
  that hold them in place

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Energetics of Melting and Freezing

• when a solid melts, it absorbs heat from its
  surroundings, it is endothermic
• as heat flows out of the surroundings, it causes
  the surroundings to cool
• as ice in your drink melts, it cause the liquid
  to cool
• when a liquid freezes, it releases heat into its
  surroundings, it is exothermic
• as heat flows into the surroundings, it causes
  the surroundings to warm

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Heat of Fusion

• the amount of heat needed to melt one mole of a
  solid is called the heat of fusion
• DHfus
• fusion is an old term for heating a substance
  until it melts, it is not the same as nuclear
  fusion
• since freezing is the opposite process to
  melting, the same amount of energy transferred is
  the same, but in the opposite direction
• DHcrystal -DHfus
• in general, DHvap gt DHfus because vaporization
  requires breaking all attractive forces

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Heats of Fusion of Several Substances
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Sublimation

• sublimation is a physical change in which the
  solid form changes directly to the gaseous form
• without going through the liquid form
• like melting, sublimation is endothermic

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IntermolecularAttractive Forces
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Why are molecules attracted to each other?

• intermolecular attractions are due to attractive
  forces between opposite charges
• ion to - ion
• end of polar molecule to - end of polar
  molecule
• H-bonding especially strong
• larger charge stronger attraction
• even nonpolar molecules will have a temporary
  induced dipoles

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Dispersion Forces

• also known as London Forces or Induced Dipoles
• caused by electrons on one molecule distorting
  the electron cloud on another
• all molecules have dispersion forces

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Instantaneous Dipoles
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Strength of the Dispersion Force

• depends on how easily the electrons can move, or
  be polarized
• the more electrons and the farther they are from
  the nuclei, the larger the dipole that can be
  induced
• strength of the dispersion force gets larger with
  larger molecules

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Attractive Forces and Properties

• stronger attractive forces between molecules
  higher boiling point
• in pure substance
• stronger attractive forces between molecules
  higher melting point
• in pure substance
• though also depends on crystal packing

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Dispersion Force and Molar Mass
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Permanent Dipoles

• because of the kinds of atoms that are bonded
  together and their relative positions in the
  molecule, some molecules have a permanent dipole
• all polar molecules have a permanent dipole

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Dipole-to-Dipole Attraction

• polar molecules have a permanent dipole
• a end and a end
• the end of one molecule will be attracted to
  the end of another

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Polarity and Dipole-to-Dipole Attraction
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Attractive Forces
Dispersion Forces all molecules
Dipole-to-Dipole Forces polar molecules
-
-
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Attractive Forces and Properties

• Like dissolves Like
• miscible liquids that do not separate, no
  matter what the proportions
• polar molecules dissolve in polar solvents
• water, alcohol, CH2Cl2
• molecules with O or N higher solubility in H2O
  due to H-bonding with H2O
• nonpolar molecules dissolve in nonpolar solvents
• ligroin (hexane), toluene, CCl4
• if molecule has both polar nonpolar parts, then
  hydrophilic - hydrophobic competition

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Immiscible Liquids
When liquid pentane, a nonpolar substance, is
mixed with water, a polar substance, the two
liquids separate because they are more attracted
to their own kind of molecule than to the other.
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Hydrogen Bonding

• Molecules that have HF, OH or NH groups have
  particularly strong intermolecular attractions
• unusually high melting and boiling points
• unusually high solubility in water
• this kind of attraction is called a Hydrogen Bond

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Properties and H-Bonding
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Intermolecular H-Bonding
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Hydrogen Bonding

• When a very electronegative atom is bonded to
  hydrogen, it strongly pulls the bonding electrons
  toward it.
• Since hydrogen has no other electrons, when it
  loses the electrons, the nucleus becomes
  deshielded
• exposing the proton
• The exposed proton acts as a very strong center
  of positive charge, attracting all the electron
  clouds from neighboring molecules

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H-Bonds vs. Chemical Bonds

• hydrogen bonds are not chemical bonds
• hydrogen bonds are attractive forces between
  molecules
• chemical bonds are attractive forces that make
  molecules

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Attractive Forces Properties
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Types of Intermolecular Forces
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Crystalline Solids
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Types of Crystalline Solids
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Molecular Crystalline Solids

• Molecular solids are solids whose composite units
  are molecules
• Solid held together by intermolecular attractive
  forces
• dispersion, dipole-dipole, or H-bonding
• generally low melting points and DHfusion

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Ionic Crystalline Solids

• Ionic solids are solids whose composite units are
  formula units
• Solid held together by electrostatic attractive
  forces between cations and anions
• cations and anions arranged in a geometric
  pattern called a crystal lattice to maximize
  attractions
• generally higher melting points and DHfusion than
  molecular solids
• because ionic bonds are stronger than
  intermolecular forces

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Atomic Crystalline Solids

• Atomic solids are solids whose composite units
  are individual atoms
• Solid held together by either covalent bonds,
  dispersion forces or metallic bonds
• melting points and DHfusion vary depending on the
  attractive forces between the atoms

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Types of Atomic Solids
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Types of Atomic SolidsCovalent

• Covalent Atomic Solids have their atoms attached
  by covalent bonds
• effectively, the entire solid is one, giant
  molecule
• because covalent bonds are strong, these solids
  have very high melting points and DHfusion
• because covalent bonds are directional, these
  substances tend to be very hard

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Types of Atomic SolidsNonbonding

• Nonbonding Atomic Solids are held together by
  dispersion forces
• because dispersion forces are relatively weak,
  these solids have very low melting points and
  DHfusion

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Types of Atomic SolidsMetallic

• Metallic solids are held together by metallic
  bonds
• metal atoms release some of their electrons to be
  shared by all the other atoms in the crystal
• the metallic bond is the attraction of the metal
  cations for the mobile electrons
• often described as islands of cations in a sea of
  electrons

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Metallic Bonding

• the model of metallic bonding can be used to
  explain the properties of metals
• the luster, malleability, ductility, electrical
  and thermal conductivity are all related to the
  mobility of the electrons in the solid
• the strength of the metallic bond varies,
  depending on the charge and size of the cations
  so the melting points and DHfusion of metals vary
  as well

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Water A Unique and Important Substance

• water is found in all 3 states on the earth
• as a liquid, it is the most common solvent found
  in nature
• without water, life as we know it could not exist
• the search for extraterrestrial life starts with
  the search for water

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Water

• liquid at room temperature
• most molecular substances that have a molar mass
  (18.02 g/mol) similar to waters are gaseous
• relatively high boiling point
• expands as it freezes
• most substances contract as they freeze
• causes ice to be less dense than liquid water