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Chapter 12 Liquids, Solids, and Intermolecular Forces. Introductory Chemistry , 3 rd Edition Nivaldo Tro. Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA. 2009, Prentice Hall. Interactions Between Molecules.

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introductory chemistry 3 rd edition nivaldo tro

Chapter 12

Liquids, Solids, and Intermolecular Forces

Introductory Chemistry, 3rd EditionNivaldo Tro

Roy Kennedy

Massachusetts Bay Community College

Wellesley Hills, MA

2009, Prentice Hall

interactions between molecules
Interactions Between Molecules
  • Many of the phenomena we observe are related to interactions between molecules that do not involve a chemical reaction, such as taste, smell at the receptor sites in our tongue and nose, fleeting in nature.

Tro\'s Introductory Chemistry, Chapter 12

interactions between molecules3
Interactions Between Molecules
  • Important to scientists are the physical properties of solids, liquids and gas:
    • How can water exists in three states on this planet
    • Why are there solids
    • What holds a liquid together
    • Why are gasses not liquids
    • Why does matter have such different melting and boiling points

Tro\'s Introductory Chemistry, Chapter 12

the physical states of matter
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.

Tro\'s Introductory Chemistry, Chapter 12

structure determines properties
Structure Determines Properties
  • The atoms or molecules have different structures in solids, liquids, and gases, leading to different properties.

Tro\'s Introductory Chemistry, Chapter 12

properties of the states of matter gases
Properties of the States of Matter:Gases
  • Low densities compared to solids and liquids.
  • Fluid.
    • The material exhibits a smooth, continuous flow as it moves.
  • Take the shape of their container(s).
  • Expand to fill their container(s).
  • Can be compressed into a smaller volume.

Tro\'s Introductory Chemistry, Chapter 12

properties of the states of matter liquids
Properties of the States of Matter:Liquids
  • High densities compared to gases.
  • Fluid.
    • The material exhibits a smooth, continuous flow as it moves.
  • Take the shape of their container(s).
  • Keep their volume, do not expand to fill their container(s).
  • Cannot be compressed into a smaller volume.

Tro\'s Introductory Chemistry, Chapter 12

properties of the states of matter solids
Properties of the States of Matter:Solids
  • High densities compared to gases.
  • Nonfluid.
    • They move as an entire “block” rather than a smooth, continuous flow.
  • Keep their own shape, do not take the shape of their container(s).
  • Keep their own volume, do not expand to fill their container(s).
  • Cannot be compressed into a smaller volume.

Tro\'s Introductory Chemistry, Chapter 12

the structure of solids liquids and gases
The Structure of Solids, Liquids, and Gases

Tro\'s Introductory Chemistry, Chapter 12

gases
Gases
  • In the gas state, the particles have complete freedom from each other.
  • The particles are constantly flying around, bumping into each other and their container(s).
  • In the gas state, there is a lot of empty space between the particles.
    • On average.

Tro\'s Introductory Chemistry, Chapter 12

gases continued
Gases, Continued
  • 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(s), and will flow.

Tro\'s Introductory Chemistry, Chapter 12

liquids
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 don’t have enough freedom to escape and expand to fill the container(s).

Tro\'s Introductory Chemistry, Chapter 12

solids
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.

Tro\'s Introductory Chemistry, Chapter 12

solids continued
Solids, Continued
  • 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.

Tro\'s Introductory Chemistry, Chapter 12

why is sugar a solid but water is a liquid
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.
    • They depend on shape, polarity, etc.
  • The ability of the molecules to overcome the attraction depends on the amount of kinetic energy they possess.

Tro\'s Introductory Chemistry, Chapter 12

properties and attractive forces
Properties and Attractive Forces

Tro\'s Introductory Chemistry, Chapter 12

phase changes melting
Phase Changes:Melting
  • Generally, we convert a material in the solid state into a liquid by heating it.
  • Adding heat energy increases the amount of kinetic energy of the molecules in the solid.
  • Eventually, they acquire enough energy to partially overcome the attractive forces holding them in place.
  • This allows the molecules enough extra freedom to move around a little and rotate.

Tro\'s Introductory Chemistry, Chapter 12

phase changes boiling
Phase Changes:Boiling
  • Generally, we convert a material in the liquid state into a gas by heating it.
  • Adding heat energy increases the amount of kinetic energy of the molecules in the liquid.
  • Eventually, they acquire enough energy to completely overcome the attractive forces holding them together.
  • This allows the molecules complete freedom to move around and rotate.

Tro\'s Introductory Chemistry, Chapter 12

properties of liquids surface tension
Properties of Liquids:Surface Tension
  • Liquids tend to minimize their surface—a phenomenon we call surface tension.
  • This tendency causes liquids to have a surface that resists penetration.
  • The stronger the attractive force between the molecules, the larger the surface tension.

Tro\'s Introductory Chemistry, Chapter 12

surface tension
Surface Tension
  • Molecules in the interior of a liquid experience attractions to surrounding molecules in all directions.
  • However, 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.

Tro\'s Introductory Chemistry, Chapter 12

properties of liquids viscosity
Properties of Liquids:Viscosity
  • Some liquids flow more easily than others.
  • The resistance of a liquid’s flow is called viscosity.
  • The stronger the attractive forces between the molecules, the more viscous the liquid is.
  • Also, the less round the molecule’s shape, the larger the liquid’s viscosity.
    • Some liquids are more viscous because their molecules are long and get tangled in each other, causing them to resist flowing.

Tro\'s Introductory Chemistry, Chapter 12

evaporation
Evaporation
  • The evaporation happens at the surface.
  • Molecules on the surface experience a smaller net attractive force than molecules in the interior.
  • All the surface molecules do not escape at once, only the ones with sufficient kinetic energy to overcome the attractions will escape.

Tro\'s Introductory Chemistry, Chapter 12

escaping the surface
Escaping the Surface
  • The average kinetic energy is directly proportional to the Kelvin temperature.
  • Not all molecules in the sample have the same amount of 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.
escaping the surface continued
Escaping the Surface, Continued
  • 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 that the liquid continues to evaporate .

Tro\'s Introductory Chemistry, Chapter 12

factors effecting the rate of evaporation
Factors Effecting the Rate of Evaporation
  • Liquids that evaporate quickly are called volatile liquids, while those that do not are called nonvolatile.
  • Increasing the surface area increases the rate of evaporation.
    • More surface molecules.
  • Increasing the temperature increases the rate of evaporation.
    • Raises the average kinetic energy, resulting in more molecules that can escape.
  • Weaker attractive forces between the molecules = faster rate of evaporation.

Tro\'s Introductory Chemistry, Chapter 12

reconnecting with the surface
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.

Tro\'s Introductory Chemistry, Chapter 12

evaporation and condensation
Evaporation and Condensation

Shortly, the water

starts to evaporate.

Initially the rate

of evaporation is much faster than rate of condensation

When water is just

added to the flask and it is capped, all the water molecules are in the liquid.

Opposite processes that occur at the same rate in the same system are said to be in dynamic equilibrium.

Eventually, the condensation and evaporation reach the same speed. The air in the flask is now saturated with water vapor.

Tro\'s Introductory Chemistry, Chapter 12

vapor pressure
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 don’t 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.

Tro\'s Introductory Chemistry, Chapter 12

boiling
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.

Tro\'s Introductory Chemistry, Chapter 12

boiling point
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.

Tro\'s Introductory Chemistry, Chapter 12

temperature and boiling
Temperature and Boiling
  • As you heat a liquid, its temperature increases until it reaches its boiling point.
  • All the energy from the heat source is being used to overcome all of the attractive forces in the liquid.

Tro\'s Introductory Chemistry, Chapter 12

energetics of evaporation
Energetics of Evaporation
  • As it loses its 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.

Tro\'s Introductory Chemistry, Chapter 12

energetics of condensation
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.

Tro\'s Introductory Chemistry, Chapter 12

heat of vaporization
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
    • 100 °C.
    • Always endothermic.
      • Number is +.

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.
    • DHcondensation = −DHvaporization

Tro\'s Introductory Chemistry, Chapter 12

heats of vaporization of liquids at their boiling points and at 25 c
Heats of Vaporization of Liquidsat Their Boiling Points and at 25 °C

Tro\'s Introductory Chemistry, Chapter 12

example 12 1 calculate the mass of water that can be vaporized with 155 kj of heat at 100 c

kJ

mol H2O

g H2O

Example 12.1—Calculate the Mass of Water that Can Be Vaporized with 155 KJ of Heat at 100 °C.

Given:

Find:

155 kJ

g H2O

Solution Map:

Relationships:

1 mol H2O = 40.7 kJ, 1 mol = 18.02 g

Solution:

Check:

Since the given amount of heat is almost 4x the DHvap, the amount of water makes sense.

slide37
Practice—How Much Heat Energy, in kJ, is Required to Vaporize 87 g of Acetone, C3H6O, (MM 58.08) at 25 C? (DHvap = 31.0 kJ/mol)

Tro\'s Introductory Chemistry, Chapter 12

slide38

g C3H6O

mol C3H6O

kJ

Practice—How Much Heat Energy, in kJ, Is Required to Vaporize 87 g of Acetone, C3H6O, (MM 58.08) at 25 C? (DHvap = 31.0 kJ/mol), Continued

Given:

Find:

87 g C3H6O

kJ

Solution Map:

Relationships:

1 mol C3H6O = 31.0 kJ at 25 C, 1 mol = 58.08 g

Solution:

Check:

Since the given mass is than one mole, the answer being greater than DHvap makes sense.

temperature and melting
Temperature and Melting
  • As you heat a solid, its temperature increases until it reaches its 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 some of the attractive forces in the solid that hold them in place.

Tro\'s Introductory Chemistry, Chapter 12

energetics of melting and freezing
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 heat flows out of your drink into the ice cubes (causing them to melt), the liquid gets cooler.
  • 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.
    • Orange growers often spray their oranges with water when a freeze is expected. Why?

Tro\'s Introductory Chemistry, Chapter 12

heat of fusion
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 (crystallization) is the opposite process of melting, the amount of energy transferred is the same, but in the opposite direction.
    • DHcrystal= -DHfus
  • In general, DHvap > DHfus because vaporization requires breaking all attractive forces.

Tro\'s Introductory Chemistry, Chapter 12

heats of fusion of several substances
Heats of Fusion of Several Substances

Tro\'s Introductory Chemistry, Chapter 12

slide43
Practice—How Much Heat Energy, in kJ, is Required to Melt 87 g of Acetone, C3H6O, (MM 58.08)?(DHfus = 5.69 kJ/mol)

Tro\'s Introductory Chemistry, Chapter 12

slide44

g C3H6O

mol C3H6O

kJ

Practice—How Much Heat Energy, in kJ, Is Required to Melt 87 g of Acetone, C3H6O, (MM 58.08)?, Continued

Given:

Find:

87 g C3H6O

kJ

Solution Map:

Relationships:

1 mol C3H6O = 5.69 kJ at -94.8 C, 1 mol = 58.08 g

Solution:

Check:

Since the given mass is more than one mole, the answer being greater than DHvap makes sense.

sublimation
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.

Tro\'s Introductory Chemistry, Chapter 12

intermolecular attractive forces

IntermolecularAttractive Forces

Tro\'s Introductory Chemistry, Chapter 12

effect of the strength of intermolecular attractions on properties
Effect of the Strength of Intermolecular Attractions on Properties
  • The stronger the intermolecular attractions are, the more energy it takes to separate the molecules.
  • Substances with strong intermolecular attractions have higher boiling points, melting points, and heat of vaporization; they also have lower vapor pressures.

Tro\'s Introductory Chemistry, Chapter 12

practice pick the substance in each pair with the stronger intermolecular attractions
Practice—Pick the Substance in Each Pair with the Stronger Intermolecular Attractions.
  • sugar or water
  • water or acetone
  • ice or dry ice
  • sugar or water.
  • water or acetone.
  • ice or dry ice.

Tro\'s Introductory Chemistry, Chapter 12

attractive forces and properties
Attractive Forces and Properties
  • Like dissolves like.
    • Miscible = Liquids that do not separate, no matter what the proportions.
  • Ionic
  • Polar molecules dissolve in polar solvents.
    • Water, alcohol, CH2Cl2.
    • A special case of polarity: Molecules with O or N higher solubility in H2O due to H-bonding with H2O.
  • Nonpolar molecules dissolve in nonpolar solvents or Dispersion forces
    • Ligroin (hexane), toluene, CCl4.
  • If molecule has both polar and nonpolar parts, then hydrophilic-hydrophobic competition.
dispersion forces
Dispersion Forces
  • Also known as London forces or instantaneous dipoles.
  • Caused by distortions in the electron cloud of one molecule inducing distortion in the electron cloud on another.
  • Distortions in the electron cloud lead to a temporary dipole.
  • The temporary dipoles lead to attractions between molecules—dispersion forces.
  • All molecules have attractions caused by dispersion forces.

Tro\'s Introductory Chemistry, Chapter 12

instantaneous dipoles
Instantaneous Dipoles

Tro\'s Introductory Chemistry, Chapter 12

strength of the dispersion force
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.

Tro\'s Introductory Chemistry, Chapter 12

dispersion force and molar mass
Dispersion Force and Molar Mass

Tro\'s Introductory Chemistry, Chapter 12

slide54
Practice—The Following Are All Made of Non–Polar Molecules. Pick the Substance in Each Pair with the Highest Boiling Point.
  • CH4 or C3H8.
  • BF3 or BCl3.
  • CO2 or CS2.
  • CH4 or C3H8
  • BF3 or BCl3
  • CO2 or CS2

Tro\'s Introductory Chemistry, Chapter 12

permanent dipoles
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.
    • Polar molecules.
  • The size of the molecule’s dipole is measured in debyes, D.

Tro\'s Introductory Chemistry, Chapter 12

dipole to dipole attraction
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.

Tro\'s Introductory Chemistry, Chapter 12

polarity and dipole to dipole attraction
Polarity and Dipole-to-Dipole Attraction

Tro\'s Introductory Chemistry, Chapter 12

attractive forces

+ - + -

_

+

+

_

+

_

+

_

Attractive Forces

Dispersion forces—All molecules.

Dipole-to-dipole forces—Polar molecules.

+ -

+ -

Tro\'s Introductory Chemistry, Chapter 12

intermolecular attraction and properties
Intermolecular Attraction and Properties
  • All molecules are attracted by dispersion forces.
  • Polar molecules are also attracted by dipole-dipole attractions.
  • Therefore, the strength of attraction is stronger between polar molecules than between nonpolar molecules of the same size.

Tro\'s Introductory Chemistry, Chapter 12

slide60









Practice—Determine Which of the Following Has Dipole–Dipole Attractive Forces.(EN C= 2.5, F = 4, H = 2.1, S = 2.5)
  • CS2Nonpolar bonds = nonpolar molecule.
  • CH2F2Polar bonds and asymmetrical = polar molecule.
  • CF4Polar bonds and symmetrical shape = nonpolar molecule.
  • CS2
  • CH2F2
  • CF4

Tro\'s Introductory Chemistry, Chapter 12

attractive forces and properties61
Attractive Forces and Properties
  • Like dissolves like.
    • Miscible = Liquids that do not separate, no matter what the proportions.
  • Ionic
  • 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 and nonpolar parts, then hydrophilic-hydrophobic competition.
immiscible liquids
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.

Tro\'s Introductory Chemistry, Chapter 12

hydrogen bonding
Hydrogen Bonding
  • HF, or molecules that have 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.

Tro\'s Introductory Chemistry, Chapter 12

properties and h bonding
Properties and H-Bonding

Tro\'s Introductory Chemistry, Chapter 12

intermolecular h bonding
Intermolecular H-Bonding

Tro\'s Introductory Chemistry, Chapter 12

hydrogen bonding67
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.

Tro\'s Introductory Chemistry, Chapter 12

h bonds vs chemical bonds
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.

Tro\'s Introductory Chemistry, Chapter 12

attractive forces and properties69
Attractive Forces and Properties

Tro\'s Introductory Chemistry, Chapter 12

example 12 5 which of the following is a liquid at room temperature the other two are gases
Example 12.5—Which of the Following Is a Liquid at Room Temperature? (The Other Two Are Gases.)
  • formaldehyde, CH2O.
    • 30.03 g/mol.
    • polar molecule  dipole–dipole attractions present.
      • Polar C=O bond and asymmetric.
  • fluoromethane, CH3F.
    • 34.03 g/mol.
    • polar molecule  dipole–dipole attractions present.
      • Polar C−F bond and asymmetric.
  • hydrogen peroxide, H2O2
    • 34.02 g/mol.
    • polar molecule  dipole–dipole attractions present.
      • Polar H−O bonds and asymmetric.
    • H−O bonds  hydrogen-bonding present.
  • formaldehyde, CH2O
    • 30.03 g/mol
    • polar molecule  dipole–dipole attractions present
      • polar C=O bond & asymmetric
  • fluoromethane, CH3F
    • 34.03 g/mol
    • polar molecule  dipole–dipole attractions present
      • polar C−F bond & asymmetric
  • hydrogen peroxide, H2O2
    • 34.02 g/mol
    • polar molecule  dipole–dipole attractions present
      • polar H−O bonds & asymmetric
    • H−O bonds  Hydrogen bonding present

Tro\'s Introductory Chemistry, Chapter 12

practice pick the compound in each pair expected to have the higher solubility in h 2 o
Practice–Pick the Compound in Each Pair Expected to Have the Higher Solubility in H2O.
  • CH3CH2OCH2CH3 or CH3CH2CH2CH2CH3.
  • CH3CH2NHCH3 or CH3CH2CH2CH3.
  • CH3CH2OH or CH3CH2CH2CH2CH2OH.

Tro\'s Introductory Chemistry, Chapter 12

practice pick the compound in each pair expected to have the higher solubility in h 2 o continued
Practice–Pick the Compound in Each Pair Expected to Have the Higher Solubility in H2O, Continued.
  • CH3CH2OCH2CH3 or CH3CH2CH2CH2CH3 contains polar O.
  • CH3CH2NHCH3 or CH3CH2CH2CH3contains polar N.
  • CH3CH2OH or CH3CH2CH2CH2CH2OH contains less nonpolar parts.

Tro\'s Introductory Chemistry, Chapter 12

crystalline solids

Crystalline Solids

Tro\'s Introductory Chemistry, Chapter 12

types of crystalline solids
Types of Crystalline Solids

Tro\'s Introductory Chemistry, Chapter 12

molecular crystalline solids
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.

Tro\'s Introductory Chemistry, Chapter 12

ionic crystalline solids
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.

Tro\'s Introductory Chemistry, Chapter 12

atomic crystalline solids
Atomic Crystalline Solids
  • Atomic solids are solids whose composite units are individual atoms.
  • Solids held together by either covalent bonds, dispersion forces, or metallic bonds.
  • Melting points and DHfusion vary depending on the attractive forces between the atoms.

Tro\'s Introductory Chemistry, Chapter 12

practice classify each of the following crystalline solids as molecular ionic or atomic
Practice—Classify Each of the Following Crystalline Solids as Molecular, Ionic, or Atomic.
  • H2O(s)
  • Si(s)
  • C12H22O11(s)
  • CaF2(s)
  • Sc(NO3)3(s)
  • H2O(s)—molecular.
  • Si(s)—atomic.
  • C12H22O11(s)—molecular.
  • CaF2(s)—ionic.
  • Sc(NO3)3(s)—ionic.

Tro\'s Introductory Chemistry, Chapter 12

metallic bonding
Metallic Bonding
  • The model of metallic bonding can be used to explain the properties of metals.
  • The luster, malleability, ductility, and 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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