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Topic 11

Liquids and Solids

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States of MatterComparison of gases, liquids, and solids. – Gases are compressible fluids. Their molecules are

widely separated with no volume or shape.– Liquids are relatively incompressible fluids. Their

molecules are more tightly packed and touching with volume but no shape.

– Solids are nearly incompressible and rigid. Their molecules or ions are in close contact and do not move with volume and shape.

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Changes of StateA change of state or phase transition is a change of a substance from one state to another.

The energy associated with changing states is equal for both directions but opposite in sign meaning exothermic vs endothermic.

solid

liquid

gas

melting orfusion

freezing or crystallization

condensationboiling or vaporization

sublimation condensation or deposition

exo

exo

exo

endo

endo

endo

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Vapor PressureLiquids are continuously vaporizing even if their not at their boiling point.

– If a liquid is in a closed vessel with space above it, a partial pressure of the vapor state builds up in this space.

– The vapor pressure of a liquid is the partial pressure of the vapor over the liquid, measured at equilibrium at a given temperature.

– The more gas molecules (vapor), the higher the vapor pressure.

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Vapor PressureThe vapor pressure of a liquid depends on its temperature.

– Liquids with relatively high vapor pressures at normal temperatures are said to be volatile. The higher the vapor pressure, the easier it goes from a liquid to a gas.

– i.e. acetone has a high vapor pressure.

– As the temperature increases, the kinetic energy of the molecular motion becomes greater, and vapor pressure increases. More energy available to convert liquid molecules to gas; hence, more gas molecules and higher vapor pressure.

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Boiling PointThe temperature at which the vapor pressure of a liquid equals the pressure exerted on the liquid is called the boiling point.– As the temperature of a liquid increases, the vapor

pressure increases (more gas molecules formed) until it reaches atmospheric pressure and begins to boil.

– At this point, stable bubbles of vapor (same species) form within the liquid. This is called boiling.

– The normal boiling point is the boiling point at 1 atm.– Because atmospheric pressure varies with altitude and

weather conditions, the boiling point of a liquid does as well.

– Boiling point of water at 1 atm is 100oC but approximately 71oC on Mount Everest at 8850m elevation (0.33 atm).

as atm pressure decreases, boiling point decreases

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Clausius-Clapeyron EquationWe noted that vapor pressure was a function of temperature.

– It has been demonstrated that the logarithm of the vapor pressure of a liquid varies linearly with absolute temperature, K.

)(21

vap

1

2

T1

T1

RH

PP ln

– Consequently, the vapor pressure of a liquid at two different temperatures is described by:

– Equation allows you to determine the vapor pressure of substance at any temperature if you know the vapor pressure at the normal boiling point.

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Carbon disulfide, CS2, has a normal boiling point of 46°C (vapor pressure = 760 mmHg) and a heat of vaporization of 26.8 kJ/mol. What is the vapor pressure of carbon disulfide at 35°C?

Substituting into the Clausius-Clapeyron equation, we obtain:

Hg) mm (760

P ln 2 )(K 308

1K 319

1K)J/(mol 8.31

J/mol 10 26.8 3

361.0 )K 10 (-1.12 K) (3225 1-4-

Taking the antiln we obtain:

361)antiln(-0. Hg) mm (760

P 2

Hg mm 760 361)antiln(-0. P2 Hg mm 530 P2

319 K

308 K )(21

vap

1

2

T1

T1

RH

PP ln

Note: Temp decreased and VP decreased

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Freezing PointThe temperature at which a pure liquid changes to a crystalline solid, or freezes, is called the freezing point.

– The melting point is identical to the freezing point and is defined as the temperature at which a solid becomes a liquid.

– Unlike boiling points, melting points are not affected significantly by pressure changes; however, large pressure changes may have some affect.

S L

melting, +H endothermic

freezing, -H exothermic

kJ 01.6H );l(OH)s(OH fus22

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Heat of Phase TransitionTo melt a pure substance at its melting point requires an extra boost of energy to overcome lattice energies.

– The heat needed to melt 1 mol of a pure substance is called the heat of fusion and denoted Hfus.

– For ice, the heat of fusion is 6.01 kJ/mol.

kJlOHsOH 01.6H );()( crystal22

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Heat of Phase TransitionTo boil a pure substance at its boiling point requires an extra boost of energy to overcome intermolecular forces.

– The heat needed to boil 1 mol of a pure substance is called the heat of vaporization and denoted Hvap.

kJ 66.40H );g(OH)l(OH vap22

– For liquid water, the heat of vaporization is 40.66 kJ/mol.

kJgOHlOH 66.40H );()( cond22

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Figure : Heating curve for water.

Note: Temp does not change while changing states; heat is used to overcome forces of attraction

Hfus = 6.01 kJ Hcrystal = -6.01 kJ

Hvap = 40.66 kJ Hcond = -40.66 kJ

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Phase Diagrams

A phase diagram is a graphical way to summarize the conditions under which the different states of a substance are stable for different temperatures and pressures.

– The diagram is divided into three areas representing each state of the substance.

– The curves separating each area represent the boundaries of phase changes.

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Phase DiagramsBelow is a typical phase diagram. It consists of three curves that divide the diagram into regions labeled “solid, liquid, and gas”.

B

temperature

pres

sure

A

C

D

solid liquid

gas

.

.

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Phase Diagrams

Curve AB, dividing the solid region from the liquid region, represents the conditions under which the solid and liquid are in equilibrium.

B

temperature

pres

sure

A

C

D

solid liquid

gas

.

.

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Phase DiagramsUsually, the melting point is only slightly affected by pressure. For this reason, the melting point curve, AB, is nearly vertical.

B

temperature

pres

sure

A

C

D

solid liquid

gas

.

.

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Phase Diagrams

If a liquid is more dense than its solid (i.e. water), the curve leans slightly to the left (points toward less dense state), causing the melting point to decrease with pressure.

B

temperature

pres

sure

A

C

D

solidless dense

liquidmore dense

gas

.

.

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Phase DiagramsIf a liquid is less dense than its solid, the curve leans slightly to the right, causing the melting point to increase with pressure. Most solids are more dense than liquids.

B

temperature

pres

sure

A

C

D

liquidless dense

gas

.

.solid

more dense

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Phase DiagramsCurve AC, which divides the liquid region from the gaseous region, represents the boiling points of the liquid for various pressures. Note that pressure has a major affect on curve and the boiling point.

B

temperature

pres

sure

A

C

D

solid liquid

gas

.

.

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Phase DiagramsCurve AD, which divides the solid region from the gaseous region, represents the vapor pressures of the solid at various temperatures. An example of a solid that goes directly to a gas is dry ice, CO2.

B

temperature

pres

sure

A

C

D

solid liquid

gas

.

.

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Phase DiagramsThe curves intersect at A, the triple point, which is the temperature and pressure where three phases (solid, liquid, gas) of a substance exist in equilibrium (i.e. water occurs at 273.15K and 4.58 Torr).

B

temperature

pres

sure

A

C

D

solid liquid

gas

.

.L

solidless dense

gs

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Phase DiagramsThe temperature above which the liquid state of a substance no longer exists regardless of pressure is called the critical temperature.

B

temperature

pres

sure

A

C

D

solid liquid

gas

.

.Tcrit

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Phase DiagramsThe vapor pressure at the critical temperature is called the critical pressure. Note that curve AC ends at the critical point, C. No liquefaction is observed above the critical temperature. When the pressure of the gas above the critical temperature is increased beyond the critical pressure, we have what is called a supercritical fluid, SCF.

B

temperature

pres

sure

A

C

D

solid liquid

gas

.

.Tcrit

Pcrit

SCF

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Intermolecular ForcesIn general, properties like boiling point, melting point, viscosity, vapor pressure, and surface tension depend on the strength of the attractive forces among the molecules.

The stronger the attractive forces, the higher the boiling point, viscosity and surface tension of a liquid; the higher the melting point of a solid; and the lower the vapor pressure of a liquid.

Molecules gain more freedom of movement as a solid melts or as a liquid vaporizes. The amount of energy that they need to overcome the forces of attraction among them increases the stronger the forces are.

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Surface TensionSurface tension is the energy required to increase the surface area of a liquid by a unit amount.

A molecule within a liquid is pulled in all directions, whereas a molecule on the surface is only pulled to the interior.

As a result, there is a tendency for the surface area of the liquid to be minimized.

To increase the surface area of a liquid requires movement of molecules within the interior, where they experience stronger attractions, to the surface. This requires energy.

The stronger the intermolecular (attractive) forces between molecules, the higher the surface tension.

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ViscosityViscosity is the resistance to flow exhibited by all liquids and gases.Viscosity can be illustrated by measuring the time required for a steel ball to fall through a column of the liquid. Even without such measurements, you know that syrup has a greater viscosity than water (thicker the fluid, the more resistant to flow).The stronger the intermolecular (attractive) forces among the molecules of the gas or liquid, the more resistant they would be to flow, the higher the viscosity. You can lower the viscosity of a substance by increasing the temperature giving the system more energy to overcome the attractive forces thereby increasing flow.

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Intermolecular ForcesThe attractive forces between two molecules can be classified into two major types:

1.) van der Waals forces which are the weak attractive forces in a large number of substances (all covalent bonded). 2.) hydrogen bonding interactions which occurs in substances containing hydrogen atoms bonded to certain very electronegative atoms (O, N, & F).

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Van der Waals ForcesVan der Waals forces can be further classified into:

1.) London dispersion forces (occurs between any pair of molecules)

2.) dipole-induced dipole interaction (occurs between a nonpolar molecule and a polar molecule)

3.) dipole-dipole interaction (occurs between two polar molecules)

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London Dispersion ForcesLondon forces refers to the force of attraction that exists between any pair of molecules and is the predominant interaction among most molecules.

It is due to temporary molecular polarizations, which occur because electrons are always moving causing a distortion of the electron cloud surrounding a molecule.

The larger the molecule, the more frequently the polarization occur.

Therefore, we expect attractions to be stronger among larger molecules. This means London forces increase with molecular weight. The larger a molecule, the more easily the electron cloud can be distorted.

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London Dispersion ForcesLet’s look at N2 and O2 as pure substances.

Both are nonpolar molecules that only have London dispersion forces of attraction. Since they only have London forces, the molecule with the larger molar mass will have more polarization and a higher boiling point, viscosity, and surface tension as well as a lower vapor pressure.

N2 O2

Mm 28g/mol 32g/mol

BP -196oC -183oC

Since O2 has a higher molar mass, we expect it to have a higher boiling point than N2.

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Dipole-Dipole InteractionsDipole-dipole interactions refers to the force of attraction that exists between polar molecules.

The dipole-dipole interaction is an attractive intermolecular force resulting from the tendency of polar molecules to align themselves positive end to negative end.

There is a higher electron density in the Cl end of the polar molecule; this end, we say, is partially negative.

The H end of the polar molecule is partially positive and is attracted to the partially negative end of a nearby HCl molecule which is a dipole-dipole interaction.

H Cldd+

H Cldd+

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Dipole-Dipole Interactions

N2 NO O2

Mm 28g/mol 30g/mol 32g/mol

BP -196oC -152oC -183oC

Let’s look at N2, NO, and O2 pure substances.

Both N2 and O2 are nonpolar molecules that only have London dispersion forces of attraction. However, NO is a polar molecule which has dipole-dipole interactions in addition to London forces to overcome which affects its boiling point, viscosity, surface tension, and vapor pressure.

The boiling point of NO will be the highest among these substances because of the additional dipole-dipole interactions in the polar molecule.

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Dipole-induced Dipole Interactions

Dipole-induced dipole interaction refers to the force of attraction that exists between a polar molecule and a nonpolar molecule.

When a nonpolar molecule comes close to the positive end of a polar molecule, its electrons would be attracted toward the polar molecule causing temporary (or induced) polarization.

Similarly, when a molecule comes close to the negative end of a polar molecule, its electrons would be repelled, again causing temporary (or induced) polarization.

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Hydrogen BondingHydrogen bonding is a force that exists between a hydrogen atom covalently bonded to a very electronegative atom (O, N, F).

– To exhibit hydrogen bonding, one of the following three structures must be present.

H N OH FH

:::

– Only N, O, and F are electronegative enough to leave the hydrogen nucleus almost stripped bare of electrons making it strongly attracted to a lone pair of a highly electronegative atom (O, N, F) in a nearby molecule.

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Hydrogen Bonding in H2O

H HO

: :

H HO

: :

H HO

: :H H

O: :

The H atoms in water are bonded to a highly electronegative O atom.

Because H is almost stripped bare of electrons, it is strongly attracted to the lone pair of the O atom in the neighboring water molecule

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Hydrogen BondingHydrogen bonding interaction is much stronger than dipole-dipole interaction and for small molecules stronger than London forces.

Hydrogen bonding accounts for the unusually high boiling point of water. Water molecules are small and a liquid at room temperature; substances made of molecules of comparable size are gaseous at room temperature.

O2 is a gas at room temperature, while H2O is a liquid even though H2O molecules are smaller than O2 molecules.

Water is a polar molecule and capable of extensive hydrogen bonding thereby raising its boiling point considerably.

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Hydrogen BondingMolecules exhibiting hydrogen bonding have abnormally high boiling points compared to molecules with similar van der Waals forces.

Which of the following are capable of exhibiting hydrogen bonding?

N2 HI HF (CH3)2O CH3OHNH3

CH4 C6H5OH H2SWithin the Lewis structure of the molecule, H must be attached to O, N, or F for hydrogen bonded to occur.

-68oC38

H2O H2S H2Se H2TeMm 18.02 34.08 80.98 129.63g/molBP -60.33oC -41.3oC -2oC100oC

Let’s look at H2O, H2S, H2Se, and H2Te as pure substances.

If you draw the Lewis structures and examine the VSEPR geometry of these molecules, you would determine that all of these molecules are polar with a same bent geometry.

This means that all of the molecules have London dispersion and dipole-dipole interactions. Since they are all polar molecules, their London dispersion forces will dictate their boiling points which will vary based on their molar mass (larger molar mass, stronger London forces).

Based on molar mass, we would predict H2Te to have the highest boiling point because it has the largest molar mass. If we extrapolate the boiling point of H2O based on the other actual boiling points, the boiling point of H2O should be -68oC.

However, since H2O has very strong hydrogen bonding to overcome as well, it’s boiling point is actually extremely high, 100oC, as compared to the other substances despite it’s low molar mass.

due to hydrogen bonding

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For which of the following pairs of molecules do we expect London dispersion forces, dipole-dipole, dipole-induced, and hydrogen bonding?

First, we must realize through Lewis structures and VSEPR that CO2, CH4 are nonpolar while H2O, HCl, NH3 are polar.

CH4 and CH4

H2O and H2O

H2O and CO2

NH3 and NH3

HCl and HCl

London

dipole-dipole

dipole-induced

hydrogen bondingLondon

London

London

London

dipole-dipole

all have polar-polar polar-nonpolar

hydrogen bonding

H–O, N, F

dipole-dipole

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Which species has the higher boiling point CS2 or CCl4?

If you draw the Lewis structures and examine the VSEPR geometry of these molecules, you would determine that both of these molecules are nonpolar with CS2 having a linear geometry and CCl4 having a tetrahedral geometry.

Since both are nonpolar molecules, they only have London dispersion forces of attraction which means the molecule with the larger molar mass (CCl4) will have more polarization and a higher boiling point.

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Which species has the higher boiling point H2O or CO?

If you draw the Lewis structures and examine the VSEPR geometry of these molecules, you would determine that H2O is a polar molecule with a bent geometry and CO is a polar molecule with a linear geometry.

Since both are polar molecules, they have London dispersion and dipole-dipole interactions. Usually in this instance, the species with the larger molar mass would have the stronger London forces and higher boiling point; however, H2O also has hydrogen bonding causing the boiling point to be much higher despite the lower molar mass.

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Intermolecular ForcesIn summary, intermolecular forces play a large role in many of the physical properties of liquids and gases. These include:

– vapor pressureas intermolecular forces increase, vapor pressure

decreases– boiling point

as intermolecular forces increase, boiling point increases– surface tension

as intermolecular forces increase, surface tension increases– viscosity

as intermolecular forces increase, viscosity decreases

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Crystalline SolidsThe regular arrangement of particles in a crystalline solid leads to the minimization of total potential energy of interactions of the particles and the most stable arrangement.

A crystalline structure is said to have a long-range order. The overall structure can be thought of in terms of a repeating pattern, called a unit cell. The unit cells making up the solid are in close contact and in fixed positions.

Solids are characterized by the type of force holding the structural units together. In some cases, these forces are intermolecular, but in others they are chemical bonds (metallic, ionic, or covalent).

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Crystalline SolidProperties of crystals depend on the type of particles in the lattice. Crystals can be classified as ionic, molecular, or atomic.

– Ionic crystals tend to have very high melting points due to strong attractions among ions. They are brittle due to the strong repulsions that result when ions of like charges are momentarily brought closer together as ions are slightly displaced from their locations when the crystal is, say, hit by a hammer.

– Molecular crystals tend to have low melting or sublimation points. The attractive forces among the molecules are relatively weak (mainly van der Waals forces).

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Crystalline SolidAtomic crystals can be classified as nonbonding, metallic, or covalent network.– Nonbonding atomic crystals are formed when noble gases are

frozen to very low temperatures. These atoms are held together by very weak London dispersion forces.

– Metallic crystals are made up of atoms of metallic elements. If more than one element is present, the solid is a solution and is called an alloy. A strong metallic bond is the reason metals have high melting points and boiling points; most metals are solids at room temperature. Because metal atoms can readily slip and roll over each other without breaking the metallic bond, metals are malleable and ductile.

– Covalent network crystal can be thought of as one giant molecule; the atoms are held together by very strong covalent bonds. Thus, a network covalent crystal like diamond has a very high melting point and is among the hardest material known.

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Physical PropertiesMany physical properties of a solid can be attributed to its structure and forces of attraction called crystal lattice energy or ion-ion intermolecular forces.

– For a solid to melt, the forces holding the structural units together must be overcome.

– For a molecular solid, these are weak intermolecular attractions.

– Thus, molecular solids tend to have low melting points (below 300oC).

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Physical Properties– For ionic solids and covalent network solids

to melt, chemical bonds must be broken.– For that reason, their melting points are

relatively high.

– Note that for ionic solids, melting points increase with the strength of the ionic bond while solubility decreases.

– Ionic bonds are stronger when:1. The magnitude of charge is high. Higher the charge, the

stronger the attraction, the more energy needed to overcome attraction; therefore, MP increases and solubility decreases.

2. The ions are small (higher charge density). Smaller the radius, the closer the opposite charges and larger attraction, the more energy needed to overcome attraction; therefore, MP increases and solubility decreases.

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Summary: The attractive forces (crystal lattice energy) between a pair of oppositely charged ions increases (stronger bond) as the charges on the ions increases and as ionic size decreases; hence higher MP and lower solubility.

Which of the following has the higher melting point and lower solubility?

MgO vs. NaCl

CaBr2 vs. CaCl2

HW 69

Since magnesium oxide involves higher charges (+2, -2) than sodium chloride (+1, -1), MgO will have the higher MP and lower solubility.

Since both species have the same charges (+2, -1), the size of the ions (anion in this case) will affect the properties. Calcium chloride has the smaller anion; therefore, it will have the stronger attraction to calcium and have the higher MP and lower solubility.code: liquids