Liquids and Solids

Intermolecular Forces

• Kinetic molecular theory states that attractive forces between molecules in a gas are negligible, resulting in no definite shape or volume.

• Liquids and solids possess definite volumes due to attractive forces holding particles together.

• Temporary attractive forces are known as intermolecular forces, applicable to atoms, ions, and molecules.

• Intermolecular forces are weaker than covalent or ionic bonds but are based on electrostatic attractions.

Intermolecular vs. Intramolecular Forces

• Atoms within a molecule are held together by intramolecular covalent bonds.

• Different molecules are held together by weak, temporary intermolecular forces.

• These forces are attractions between areas of higher electron density in one molecule and lower electron density in another.

• Intermolecular forces frequently break and reform between different molecules in a liquid sample.

Types of Intermolecular Forces

• Ion–dipole attractions: Form between ions and polar molecules.

• Dipole–dipole attractions: Form between two polar molecules.

• Hydrogen bonding: A particularly strong type of dipole–dipole attraction.

• Dispersion forces: Occur between all molecules but are most notable between nonpolar molecules.

• The strength and type of intermolecular forces can explain trends in melting and boiling points and solubilities and can explain many of the unique properties of water.

Ion–Dipole Attractions

• Polar molecules, such as H_2O, have separate locations of partial positive and partial negative charges; they are dipoles.

• Ionic compounds dissolve in water to form hydrated ions.

• Hydrated ions form by way of ion–dipole attractions.

• Ion–dipole attractions are the strongest of the temporary attractions between particles.

Example 12.1: Pairs of compounds that can form ion-dipole attractions: NaCl and H2O.

Dipole–Dipole Attractions

• Polar molecules interact with other polar molecules.

• The negative pole of one polar molecule is attracted to the positive pole of another polar molecule.

• These attractions are a type of intermolecular force known as dipole–dipole attractions.

Example 12.2: Compounds that can form dipole-dipole attractions between their molecules: CH2O.

Hydrogen Bonds

• A hydrogen bond is a special case of dipole–dipole attractions and is stronger than other dipole–dipole attractions.

• Hydrogen bonding requires two very specific components:

  • A hydrogen bond donor molecule containing a partially positive H atom bonded to N, O, or F

  • A hydrogen bond acceptor molecule containing a partially negative N, O, or F with lone pair electrons

• A hydrogen bond is the attraction between the δ+ H in the donor molecule and a lone pair of electrons on the δ− O, N, or F of the acceptor molecule.

Example 12.3: Substances that can form hydrogen bonds to other identical molecules: HF.

Importance of Hydrogen Bonding

• Hydrogen bonding is responsible for the low density of ice compared with liquid water.

• Because ice is less dense than water, it floats on top of water and acts as a blanket to protect the aquatic life below from the extremes of cold that occur above the ice.

• Hydrogen bonding is involved in the helical structure of DNA and the shapes of certain proteins and other large biomolecules in plants and animals.

Dispersion Forces

• The electrons surrounding atoms are in constant random motion.

• At any instant in time, a molecule or atom may have an asymmetric distribution of electrons, forming an instantaneous dipole.

• An asymmetric charge distribution in one molecule can induce a similar charge asymmetry in an adjacent molecule, forming an induced dipole.

• The resulting short-lived attractions between instantaneous dipoles and induced dipoles are dispersion forces.

• The instantaneous dipole-induced dipole attraction may last for only an instant, making this type of intermolecular force far weaker than dipole–dipole attractions.

• All atoms and molecules exhibit dispersion forces.

• These forces are especially important for understanding the physical properties of nonpolar substances.

Factors Affecting the Strength of Dispersion Forces

• The ability of an electron cloud to become asymmetric, its polarizability, affects the strength of its dispersion forces.

• Smaller atoms do not polarize readily.

• Outer electrons are close to the nucleus.

• There are few core electrons for shielding from the nuclear charge.

• Larger atoms polarize more readily.

• Outer electrons farther from nucleus

• There are more core electrons for shielding.

Atomic Size and Polarizability

• Consider the diatomic halogens: F2, Cl2, Br2 , and I2.

• Fluorine and chlorine, the smaller halogens, are gases at room temperature.

• Bromine, which is larger, is a liquid.

• Iodine, which is still larger, is a solid.

• However, both bromine and iodine vaporize readily, indicating the weakness of the dispersion forces.

Factors Affecting the Strength of Dispersion Forces: Atomic Size: Number of Polarizable Electrons

• Molecules that can form more instantaneous dipoles at one time are able to form stronger dispersion forces.

• Consider the boiling points of the unbranched alkanes, a family of compounds containing C and H atoms connected in long chains.

• As the alkanes increase in size, their boiling points also increase.

• Why use boiling points as a measure? Molecules must overcome their intermolecular forces to escape into the vapor phase.

Factors Affecting the Strength of Dispersion Forces: Molecular Shape

• Molecules with shapes that allow them to interact with nearby molecules form stronger dispersion forces.

• Typically, molecules that have overall linear shapes form stronger dispersion forces than those that form bulky shapes.

• Compare branched and unbranched alkanes with the same molecular formula.

• Unbranched (linear) alkanes have higher boiling points than branched alkanes, indicating stronger intermolecular forces.

Example 12.5: Substance expected to have the stronger dispersion forces—C8H18.

Types of Intermolecular Forces from Strongest to Weakest

• Ion–dipole

• Hydrogen bonding

• Dipole–dipole

• Dispersion

Predicting the Predominant Intermolecular Force of a Molecule

• For small molecules, the relative strength of intermolecular forces follows a predictable order, with ion–dipole being the strongest and dispersion being the weakest.

• For large molecules, the collective dispersion forces can be extremely strong, making it tricky to rank their strength relative to other intermolecular forces.

• Ion–dipole forces can occur only in a mixture (i.e., an ionic substance mixed with a polar molecular substance).

• The other three types of intermolecular forces can occur in a pure substance or in a mixture because they can occur between identical molecules or between different molecules.

Example 12.7: Identifying all of the intermolecular forces in these molecules and specify which one predominates: PCl3.

Properties of Liquids

• The particles of a liquid are in contact with one another but not as closely as in the solid state.

• The particles of a liquid are in constant motion and interact via intermolecular forces, which are weak and temporary.

• These weak and temporary forces give liquids their fluid properties of viscosity, surface tension, and capillary action.

Viscosity

• The resistance to flow, viscosity, is determined by the strength of the intermolecular attractions and temperature.

• As temperature increases, the molecules increase in kinetic energy and move faster, overcoming some of the intermolecular attractions.

• This decreases the overall viscosity of the liquid and allows the liquid to flow more freely.

Example 12.8: Based on the Lewis structures of water and hydrogen peroxide, hydrogen peroxide has a higher viscosity at room temperature.

Surface Tension

• Another property of liquids that is affected by intermolecular forces is surface tension, the tendency of a liquid to minimize its surface.

• Molecules in the center of a liquid form intermolecular forces with other molecules of the liquid in all directions, resulting in no net pull in any direction for any molecule.

• Molecules at the surface of the liquid, however, have no liquid molecules above them, so they experience a downward pull.

Examples of Surface Tension

• Water beads up on the surface of a freshly waxed car because wax is nonpolar and cannot form strong intermolecular forces with water.

• Mercury, a large atom, is a liquid at room temperature with a very high surface tension, forming beads when spilled.

• Hexane or gasoline, with very low surface tension, spread out into a thin layer when spilled.

Cohesion and Adhesion

• Attraction between like particles is cohesion.

  • Water and mercury exhibit strong cohesion because they have strong intermolecular forces.

• Attraction between different particles is adhesion.

  • Water exhibits adhesion to other hydrogen-bonding materials.

  • Paper and cotton are both composed of cellulose, which contains many O–H bonds that can form hydrogen bonds to water.

  • Thus, paper towels and cotton kitchen towels both work very well at cleaning up spilled water.

Capillary Action

• Cohesion tells you that a liquid is a network of loosely connected molecules.

• Adhesion tells you that individual molecules of this network can be attracted to other particles, such as the inner surface of a narrow tube.

• The ability of a liquid to flow against gravity up a narrow tube is capillary action.

• The lead molecules of a liquid attach, via intermolecular forces to the inner surface of the tube, with the rest of the network molecules following.

Shape of the Meniscus

• Water has a high surface tension and is attracted to the surface of the glass or plastic in a graduated cylinder, thus forming a concave meniscus.

• Mercury has a high surface tension but no cohesion to glass or plastic and forms a convex meniscus.

• Liquids with low surface tension tend to have a flat surface in a graduated cylinder.

Example 12.9: Ranking the surface tensions of N2, HF, and HBr as liquids, from lowest to highest: N2 < HBr < HF.

Phase Changes

• The transition from the solid phase to the liquid phase is melting, or fusion. The opposite process, of liquid to solid, is called freezing.

• The process of changing a liquid to a gas is called evaporation, or vaporization. The opposite process, of gas to liquid, is called condensation.

• Solids that enter the gas state directly do so by sublimation. The opposite process, of gas directly to solid, is called deposition.

Example 12.10: The phase change described by the equation: C18H38(s) → C18H38(l). -- fusion.

Enthalpy and Phase Changes

• Vaporization (liquid to gas) requires an input of energy, while condensation (gas to liquid) releases energy.

• The energy change for the vaporization of 1 mole of a liquid is referred to as the enthalpy of vaporization, ∆Hvap.

• The terms enthalpy of fusion, ∆Hfus, and enthalpy of sublimation, ∆Hsub, refer to the energy changes associated with the melting and sublimation, respectively, of one mole of a substance.

• These enthalpy changes are sometimes referred to as heats of vaporization, fusion, and sublimation.

• Water has a ∆Hfus of 6.01 kJ/mol and a ∆Hvap of 40.7 kJ/mol.

• Melting requires sufficient energy to overcome some of the intermolecular forces, while vaporization requires sufficient energy to overcome nearly all of the intermolecular forces.

• Sublimation is a one-step process that starts and ends at the same states as the two-step process of fusion and vaporization.

  • ΔH{sub} = ΔH{fus} + ΔH_{vap}

  • ΔH_{sub} = 6.01 \frac{kJ}{ mol} + 40.7\frac{kJ}{ mol} = 46.7 \frac{kJ}{ mol}

Calculating q for Changes of State

• When enthalpy values are given in units of kJ/mol, use q = n(ΔH)

• When enthalpy values are given in units of kJ/g, use q = m(ΔH)

Example 12.11: Energy required to melt 16.4 g of ice at 0°C: 5.47 kJ.

Heating Curve

• The heating curve for a pure substance is a graph that shows how the temperature changes as the pure substance is heated.

• For water, a heating curve begins with a sample of ice at a temperature below its melting point.

• Heat energy is steadily added to it, warming the solid.

• When the melting point is reached (0°C for water), the temperature remains constant as the heat energy is being absorbed to melt the solid.

• Once the entire sample has melted, the temperature of the liquid water rises.

• When the liquid water reaches the boiling point, the temperature remains constant as heat energy is being absorbed to vaporize the liquid.

• In a closed container, the continued addition of energy causes the temperature of the trapped steam to rise.

Example 12.12: Energy required to change a 17.0 g sample of liquid water from 87.7°C to steam at 121.0°C and 1.00 atm: 42.02 kJ .

Vapor Pressure

• Substances that vaporize easily are said to be volatile, and substances that do not easily vaporize are nonvolatile.

• Volatile substances have more gas phase molecules above the surface of the liquid, creating vapor pressure.

• Vapor pressure is affected by two factors, the strength of the intermolecular forces and the temperature.

• Liquids with stronger attractive forces have lower vapor pressures than do liquids with weaker attractive forces.

• At higher temperatures, more molecules have the minimum kinetic energy needed to escape the liquid phase and enter the gas phase.

Example 12.13: substances in order from lowest vapor pressure to highest vapor pressure at room temperature: 2-propanol, acetone, methylpropane.

Heat of Vaporization

• The amount of energy needed to vaporize one mole of a substance—the enthalpy of vaporization—is also called the heat of vaporization and is always a positive value.

• Heats of vaporization are slightly temperature dependent and are generally reported at two temperatures—namely, 25°C and the normal boiling point of the substance.

Example 12.14: Energy is absorbed when 355 g of acetone is vaporized at 56.1°C: 178 kJ.

Clausius–Clapeyron Equation

• You can calculate the enthalpy of vaporization of a liquid from its vapor pressure at different temperatures using the Clausius–Clapeyron equation, shown here in a linear format:

  • lnP{vap} = -\frac{∆H{vap}}{R} (\frac{1}{T}) + lnβ

  • P_{vap} is the vapor pressure in atmospheres, R is the gas law constant (8.3145 J/mol · K), T is the temperature in kelvins, and β is a constant specific for each liquid.

• Two-Point Version of the Clausius–Clapeyron Equation

  • ln(\frac{P2}{P1}) = \frac{∆H{vap}}{R} (\frac{1}{T1} − \frac{1}{T_2})

Example 12.16: vapor pressure, in mmHg, of methanol (boiling point of 64.60°C, ΔHvap of 35.2 kJ/mol) at 25.00°C: 144 mmHg.

Boiling and Distillation

• Boiling occurs when the vapor pressure of the liquid equals the pressure of the surroundings.

• At that point, bubbles appear within the liquid itself and the vapor pressure of the liquid equals the surrounding pressure on the system.

• The normal boiling point is the boiling point of a liquid at a pressure of 1.00 atm.

• Liquids can boil at temperatures lower than their normal boiling point at lower atmospheric pressures and at higher temperatures at higher atmospheric pressures (autoclave, pressure-cooker).

Separating Volatile Substances Using Distillation

• Volatile substances can be separated from a liquid-phase mixture via distillation.

• Heat the mixture to vaporize the volatile components, direct the vapor away from the mixture, and cool it to condense it to a liquid.

• Components with higher boiling points and those that are nonvolatile are left behind in the original container.

Vapor Pressure and Dynamic Equilibrium

• In a closed system, liquid molecules escape into the gas phase and the vapor pressure builds up, resulting in condensation.

• As the vapor pressure increases, the rate of condensation increases.

• At some point, the rate of evaporation of liquid molecules will be equal to the rate of condensation of gas molecules, forming an equilibrium.

• There is no longer any net change in the vapor pressure or in the amount of liquid, even though vaporization and condensation continue to occur.

Example 12.18: What will happen to the amount of water in the liquid phase, the amount of water in the gas phase, and the vapor pressure when the sample is placed in the refrigerator: Eventually, a new equilibrium will be established at a lower Pvap.

Phase Diagrams

• Any one substance may exist in the solid, liquid, or gas phase depending on the temperature and pressure conditions.

• A phase diagram shows the phase of a specific substance under all possible pressure–temperature combinations.

• In a general phase diagram, pressure is measured on the y axis and temperature on the x axis.

• At low temperatures and high pressures, the substance is in the solid phase (upper-left region).

• At low pressures and high temperatures, the substance is in the gas phase (lower-right region).

• At moderate temperatures and pressures, the substance is a liquid (upper-middle region).

Features of Phase Diagrams: Phase Boundary Lines

• The solid lines on phase diagrams between phases represent conditions at which different phases are in equilibrium.

• The line separating solid and liquid regions indicates all the P–T conditions under which the solid and liquid states are in equilibrium.

• The line separating liquid and gas regions indicates all the P–T conditions under which the liquid and gas states are in equilibrium.

• The line separating solid and gas regions indicates all the P–T conditions under which the solid and gas states are in equilibrium.

Features of Phase Diagrams: Characteristic Points

• The single point at which these three lines intersect is called the triple point, which represents the P and T at which all three phases of the substance are in equilibrium.

• The critical point occurs at the pressure and temperature conditions above which the substance no longer exists as either a liquid or gas.

• At T and P above the critical pressure and critical temperature, the substance exists as a supercritical fluid, a fourth phase that has properties common to both liquids and gases.

• Supercritical fluids can effuse through solids as gases do but can also dissolve substances as liquids do.

Features of Phase Diagrams: Dashed Lines

• Horizontal dashed line – usually indicates 1 atm of pressure.

• Normal melting point occurs where this line crosses the solid/liquid boundary line.

• Normal boiling point occurs where this line crosses the liquid/gas boundary line.

• Other horizontal lines may be used to focus on phase transitions during temperature changes at constant pressure.

• Vertical lines can be used to focus on phase transitions during pressure changes at constant temperature.

Example 12.19: Identify the phase(s) present in regions A, B, and C in the phase diagram: gas, liquid and solid phases, respectively.

Phase Diagram for Water

• In most phase diagrams, the boundary line between the solid and liquid regions has a positive slope.

• This same boundary line in the phase diagram for water has a negative slope.

• Water has many unusual characteristics visible in a phase diagram.

• It has high melting and boiling points (and a high density) for such a small molecule.

• It is a liquid over a very broad temperature range.

• It is less dense as a solid than as a liquid.

• When pressure is applied to solid water, it melts.

Example 12.20: Describe what happens to carbon dioxide when its temperature increases from −90℃ to 25 ℃ at a pressure of 20 atm: CO2 is a solid, will melt at aproximately −50℃  to form liquid, and vaporize aproximately 0℃, and It will remain a gas until it reaches 25 ℃.

Classifying Solids by Organization of Particles

• Solids can be classified based on how their particles are organized into crystalline and amorphous solids.

  • Crystalline solids have definite melting points, abruptly forming a liquid once the melting point is reached.

    • Ice, H2O(s), and table salt are examples of crystalline solids.

  • Amorphous solids get softer as their temperature is raised, gradually forming a liquid.

    • Chocolate bars and glass are examples of amorphous solids.

Classifying Solids Based on Type of Particle

• Solids have relatively strong forces holding their particles together.

  • Molecular solids (CO2, H2O)contain molecules held to each other by intermolecular forces and have relatively low melting points.

  • Ionic solids (NaCl, MgO) are composed of ions held together by ionic bonds and have quite high melting points.

  • Covalent-network solids are composed of atoms connected by covalent bonds throughout the solid and have extremely high melting points.

  • Metallic solids are composed of metal ions loosely held together by their valence electrons and have a broad range of melting points.

Example 12.22a: he difference in their structures in simple molecular solids (such as ice) and covalent-network solids (such as diamond) are covalently bonded: In a covalent-network solid, covalent bonds connect all of the atoms in an extended network.

*Example 12.23c: the type of bonding or force holding together each of the following solids and indicate whether the melting point is low or high in ammonia, NH3: Ammonia is a molecular solid, so it is held together by intermolecular forces. *

Metallic Solids

• Metal atoms in a metallic solid neither transfer nor share electrons.

• The valence electrons in a metallic solid surround the metal cations, forming a sea of electrons.

• In the electron-sea model, each metal atom contributes it valence electrons and forms a cation with the valence electrons free to move throughout the structure.

• The electron-sea model helps to explain many metallic properties but is incomplete.

• (Band theory describes the valence electrons of a metal existing in delocalized, overlapping orbitals that extend over the solid.)

Example 12.24: Describe what an electron-sea model diagram would look like for aluminum. What charge would the ions have, and how many electrons would there be: Each atom would contribute three electrons to the sea of electron.

More About Metallic Bonding

• Metallic solids are composed of atoms in an orderly array.

• The atoms in a metal have not lost their electrons.

• When molten (liquid), metal atoms are still attracted to each other via metallic bonding, but the ions are less orderly and are mobile.

• Metallic bonding is not broken until the metal atoms vaporize, with each atom taking an appropriate number of electrons.

• Metallic bonding gives metals their unique properties.

Properties of Metals

• Metals are malleable, which means they can be hammered or bent into different shapes without breaking.

• Metals are also ductile, which means they can be drawn into long, thin wires.

• Metals are conductors of heat and electricity.

• These properties are possible because metallic bonds do not have a specific direction.

• The electrons can flow within whatever shape the metal takes.

The Unit Cell

• There are many types of crystal structures, partly because they contain different particles but also because the particles are organized into different unit cells.

• The unit cell is the simplest repeating unit of a crystal structure and arises from how the layers of particles are arranged.

• The simplest way is for the particles to align on top of one another.

Packing of Atoms in a Crystalline Structure

• How the layers are arranged in the three-dimensional structure is referred to as packing.

• Different types of packing give rise to seven basic unit cells, but this course focuses only on cubic unit cells.

• Cubic unit cells have 90° angles and edges of equal length.

• There are three types of cubic unit cells: simple cubic unit cell, body-centered cubic unit cell (bcc), and face-centered cubic unit cell (fcc).

Example 12.25: face-centered unit cells have higher density.