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Intermolecular Forces, Intermolecular Forces, Liquids, and SolidsLiquids, and Solids

Chapter 10Chapter 10

Seneca Valley SHSSeneca Valley SHS

AP ChemistryAP Chemistry

Solids, Liquids & Gases, A ComparisonSolids, Liquids & Gases, A Comparison

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The fundamental difference between states of matter is the distance between particles.

© 2009, Prentice-Hall, Inc.

Solids, Liquids & Gases, A ComparisonSolids, Liquids & Gases, A ComparisonBecause in the solid and liquid states particles are closer together, we refer to them as condensed phases.

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Intermolecular ForcesIntermolecular Forces• Chapters 8 & 9 – Bonding (AKA intramolecular

forces)• The forces holding solids and liquids together are

called intermolecular forces. • Intermolecular forces are much weaker than chemical

bonds. (< 15% as strong)• Example: energy needed to vaporize water: 40.7 kJ

compared to energy needed to break O-H bond in water: 934 kJ

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Intermolecular ForcesIntermolecular Forces• When a substance melts or boils the intermolecular

forces are broken (not the covalent bonds).• When a substance condenses intermolecular forces

are formed.Intermolecular

forces in ice

Intermolecular forces

between water molecules

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Importance of Intermolecular ForcesImportance of Intermolecular Forces

• Intermolecular forces of attraction are strong enough to control physical properties such as:• Boiling & melting points• Vapor pressures• Viscosity• Surface tension

• Collectively, the intermolecular forces being discussed are called van der Waal forces.• London dispersion• Dipole-Dipole• Hydrogen bonding

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Intermolecular Forces – Dipole-DipoleIntermolecular Forces – Dipole-Dipole

• Polar molecules are attracted to each other.• The positive end of one is attracted

to the negative end of the other and vice-versa.

• These forces are only important when the molecules are close to each other (like in solid & liquid states).

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Intermolecular Forces – Dipole-DipoleIntermolecular Forces – Dipole-Dipole

In molecules of equal mass & size, the strength of the IMF increase with increasing polarity.

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Intermolecular ForcesIntermolecular ForcesLondon Dispersion ForcesLondon Dispersion Forces• Weakest of all intermolecular forces.• Only IMF between noble gases and

nonpolar molecules.• The nucleus of one molecule (or atom)

attracts the electrons of the adjacent molecule (or atom).

• For an instant, the electron clouds become distorted.

• In that instant a dipole is formed (called an instantaneous dipole).

• Relatively weak and short-lived

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Intermolecular ForcesIntermolecular Forces

London Dispersion ForcesLondon Dispersion Forces

One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom).

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Intermolecular Forces - Intermolecular Forces - London DispersionLondon Dispersion

• Strength of dispersion forces are directly related to:• Surface area available for contact

− London dispersion forces between spherical molecules are lower than between sausage-like molecules.

• Polarizability− Polarizability - ease with which an electron cloud can be

deformed.

− Polarizability increases with the number of electrons in the molecule.

− Pi bonding also enhances the polarizability of a molecule.

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Intermolecular ForcesIntermolecular ForcesLondon Dispersion ForcesLondon Dispersion Forces

Nonpolar Molecule

Boiling Point (°C, 1 atm)

# electronsComparison of strength

H2 -253

O2 -183

Cl2 -34

Br2 59

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Intermolecular Forces – Intermolecular Forces – How to rank IMFHow to rank IMF

• Dispersion forces exist between ALL molecules.• When comparing relative strengths of

intermolecular attractions, remember:− When molecules have similar molecular

weights and shapes, dispersion forces are approximately equal. The relative strength of IMF will be measured by strengths of dipole-dipole interactions (which one is more polar).

− When molecules vary drastically in molecular weights, dispersion forces are the decisive factor in determining which substance has stronger IMF (which one is larger).

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Intermolecular ForcesIntermolecular Forces

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Intermolecular ForcesIntermolecular ForcesHydrogen BondingHydrogen Bonding• Special case of dipole-dipole forces

when H is bonded to N, O, or F.• By experiments: boiling points of

compounds with H-F, H-O, and H-N bonds are abnormally high.

• Hydrogen bonds are abnormally strong but still weaker than chemical bonds (5-25 kJ/mol versus 200-1100 kJ/mol)

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Intermolecular ForcesIntermolecular ForcesHydrogen BondingHydrogen Bonding• Hydrogen bonds are responsible for:

– Unique properties of water

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Intermolecular ForcesIntermolecular ForcesHydrogen BondingHydrogen BondingHydrogen bonds are

responsible for:– Protein Structure

• Protein folding is a consequence of H-bonding.

• DNA Transport of Genetic Information

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Intermolecular ForcesIntermolecular ForcesComparing Intermolecular ForcesComparing Intermolecular Forces

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Some Properties of LiquidsSome Properties of Liquids

ViscosityViscosity• Viscosity is the resistance of a liquid to flow.• A liquid flows by sliding molecules over each other.• The stronger the intermolecular forces, the higher the

viscosity.

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Some Properties of LiquidsSome Properties of Liquids

Surface TensionSurface Tension•Bulk molecules (those in the liquid) are equally attracted to their neighbors. Surface tension is the amount of energy required to increase the surface area of a liquid.

•Cohesive forces bind molecules to each other.

•Adhesive forces bind molecules to a surface.

•This is what leads to a meniscus!

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Some Properties of LiquidsSome Properties of Liquids

Capillary ActionCapillary Action• The rise of liquids up very narrow tubes• Helps water and dissolved nutrients move upward

through plants

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Bonding in SolidsBonding in Solids• Solids can be classified into two categories

– Amorphous – considerable disorder in their structures

– Example: Glass

– Crystalline – highly regular arrangement of atoms usually represented by a lattice

Amorphous Crystalline

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Bonding in Crystalline SolidsBonding in Crystalline Solids

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Bonding in SolidsBonding in SolidsMolecular SolidsMolecular Solids• Intermolecular forces: dipole-dipole, London

dispersion and H-bonds.• Weak intermolecular forces give rise to low melting

points.• Room temperature gases and liquids usually form

molecular solids at low temperatures.• Here is the structure of ice. Notice how the hydrogen

bonds become rigid forming a lattice linking each molecule of water together.

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Bonding in SolidsBonding in SolidsCovalent Network Solids – C & SiOCovalent Network Solids – C & SiO22

• Atoms held together by covalent bonds in large networks.

• Examples: diamond, graphite, quartz (SiO2), silicon carbide (SiC), and boron nitride (BN).

• Diamond crystal structure: Each atom here is C.

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Bonding in SolidsBonding in SolidsCovalent Network SolidsCovalent Network Solids

Crystalline SiO2 (quartz)

Each one of these structures is bonded together with regular, repeating covalent bonds.

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Bonding in SolidsBonding in SolidsIonic SolidsIonic Solids• Ions are at the lattice points.

• Electrostatic attraction keeps them frozen into this position.

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Bonding in SolidsBonding in Solids

MetalsMetals• Atoms held together by non-directional covalent

bonds in metallic crystals.

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Bonding in SolidsBonding in SolidsMetals – Electron Sea ModelMetals – Electron Sea Model• http://www.drkstreet.com/resources/metallic-bonding-animation.swf

• Atoms in a metal are arranged in a regular manner and vibrate about fixed positions

• The outermost electrons move freely, forming a ‘sea of electrons’ enveloping the positive metal ions.

• The metal ions are attracted to and hence held together by the ‘sea of electrons’ – these constititute metallic bonding.

• The movements of the electrons are random under normal conditions.

• However, when a potential difference is applied across the metal, the electrons move towards the direction of the positive pole.

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Bonding in SolidsBonding in SolidsMetals – PropertiesMetals – Properties• As a result of this non-directional bonding, metals

are:• Great conductors of heat & electricity• Ductile (able to drawn into wire)• Malleable (able to pounded into a sheet)• Able to form alloys easily

‒ Substance that contains a mixture of two elements and has metallic properties

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Bonding in SolidsBonding in SolidsMetal AlloysMetal Alloys• Two types• Substitutional – some of the metal atoms are

replaced by other metal atoms of similar size‒ Examples include brass (copper & zinc),

sterling silver (silver & copper), pewter (tin, copper, bismuth & antimony) and plumber’s solder (tin & antimony)

• Interstitial – some of the holes between the metallic crystal are filled with small atoms‒ Example would be steel (iron mixed with

carbon)

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Bonding in SolidsBonding in SolidsMetal AlloysMetal Alloys

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Phase ChangesPhase Changes• Surface molecules are only

attracted inwards towards the bulk molecules.

• Sublimation: solid gas.

• Vaporization: liquid gas.

• Melting or fusion: solid liq.

• Deposition: gas solid.

• Condensation: gas liquid.

• Freezing: liquid solid.

Energy Changes Accompanying Phase Changes

•Sublimation: Hsub > 0 (endo)

•Vaporization: Hvap > 0 (endo)

•Melting or Fusion: Hfus > 0 (endo)

•Deposition: Hdep < 0 (exo)

•Condensation: Hcon < 0 (exo)

•Freezing: Hfre < 0 (exo)

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Phase ChangesPhase Changes

Energy Changes Accompanying Phase ChangesEnergy Changes Accompanying Phase Changes• All phase changes are possible under the right

conditions (e.g. water sublimes when snow disappears without forming puddles).

• The sequence

heat solid melt heat liquid boil heat gas

is endothermic.• The sequence

cool gas condense cool liquid freeze cool solid

is exothermic.

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Phase ChangesPhase Changes

Energy Changes Accompanying Phase ChangesEnergy Changes Accompanying Phase Changes

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Phase ChangesPhase ChangesHeating CurvesHeating Curves• Plot of temperature change versus heat added is a heating curve.

• During a phase change, adding heat causes no temperature change.

– These points are used to calculate Hfus and Hvap.

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Vapor PressureVapor Pressure

Explaining Vapor Pressure on the Molecular Explaining Vapor Pressure on the Molecular LevelLevel

• Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid.

• These molecules move into the gas phase.• As the number of molecules in the gas phase

increases, some of the gas phase molecules strike the surface and return to the liquid.

• After some time the pressure of the gas will be constant at the vapor pressure.

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Vapor PressureVapor Pressure

Explaining Vapor Pressure Explaining Vapor Pressure on the Molecular Levelon the Molecular Level

• Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface.

• Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium.

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Vapor PressureVapor Pressure

Volatility, Vapor Pressure, and TemperatureVolatility, Vapor Pressure, and Temperature• If equilibrium is never established then the liquid evaporates.

• Volatile substances evaporate rapidly.

• The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates.

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Vapor PressureVapor PressureVapor Pressure and Boiling PointVapor Pressure and Boiling Point• Liquids boil when the external pressure equals the

vapor pressure.• Temperature of boiling point increases as pressure

increases.• Two ways to get a liquid to boil: increase temperature

or decrease pressure.– Pressure cookers operate at high pressure. At high

pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked.

• Normal boiling point is the boiling point at 760 mmHg (1 atm).