Liquids and solids

 

In the liquid and solid phases, particles are much closer together

·        Intermolecular attractive forces become important in determining properties

·        Liquids consist of disordered clusters of particles which can easily move randomly in three dimensions

·        Solids consist of orderly arranged particles which are largely locked into place

 

Changes of state

            Movement from a solid to a liquid (melting) or from a liquid to a gas (boiling) requires that the sample have enough energy to overcome the attractive forces holding the molecules together.

 

            A strong intermolecular force requires a high level of kinetic energy (a high temperature) to break the molecules apart.

 

 

Types of intermolecular forces:

 

Ionic – the force between neighboring fully charge ions in an ionic compound

·        Coulomb’s law: force of attraction is directly proportional to the charges on the ions and inversely proportional to the square of the distance between them      F µ (q+)(q-)/d2

·        Coulomb’s law implies that ionic compounds consisting of ions with high charges or having smaller ionic radii will have stronger forces à higher melting points

·        In the following pairs, which would have a higher melting point?

NaF                                   or                                 NaCl

KCl                                    or                                 NaCl

Na2S                                 or                                 NaCl

CaF2                                 or                                 CaO

 

 

Dipole-dipole – the force between partially charged portions of polar molecules in a polar covalent compound

·        Charges are partial and force of attraction is inversely proportional to d4, so the attraction is weaker than ion-ion attractions

 

Hydrogen bonding -- a stronger type of dipole-dipole attraction; requires that the molecules have a hydrogen attached to a small, highly electronegative atom (Nitrogen, Oxygen, or Fluorine) creating a very strong dipole due to large differences in electronegativity.

·        The small, highly electronegative element virtually removes hydrogen’s one electron – causing it to act somewhat like a bare proton

·        The d+ hydrogen atom is strongly attracted to the non-bonding electrons on the very electronegative atom in a nearby molecule

 

London forces -- the momentary force of attraction between temporarily partially charged portions of non-polar covalent molecules

·        Also called van der Waal’s forces, dispersion forces, and dipole-induced dipole forces

·        Can occur in any molecule, but is the only force available to non-polar substances and noble gases

·        Requires polarizability – a shift of electrons to one side of the electron cloud

·        Polarizability increases with increasing size/# of electrons – valence electrons are further away from the nucleus

·        London forces are stronger for molecules that are larger or have more electrons

 

Ranking of attractive forces (strongest to weakest):

1.  Ion-ion

2.  H-bonding

3.  Dipole-dipole

4.  London forces

 

 

Evaporation

            Evaporation occurs when molecules at the surface of the liquid have enough energy to overcome the attractive forces

·        Recall that kinetic energy is measured by temperature

·        At a higher temperature there will be a higher average amount of kinetic energy à more molecules will have enough energy to escape and evaporation occurs more rapidly

See Fig. 10.41 on page 485

·        Evaporation requires energy, so the remaining liquid has less energy – is cooler. 

Ex: evaporation of sweat cools the body.

·        Substances that exhibit weaker attractive forces evaporate more rapidly.  Ex. acetone vs. water

·        In a closed container, a dynamic equilibrium is reached where the rate of evaporation and the rate of condensation equalize.

·        Condensation occurs when a vapor molecule strikes a surface and attractive forces hold it in place.

 

Vapor pressure

            As a liquid evaporates, the gas particles formed emit a vapor pressure

·        A given liquid will have a higher vapor pressure at a higher temperature

·        Liquids which evaporate more quickly (due to weaker attractive forces) have a higher vapor pressure.

·        Easily vaporized liquids are referred to as volatile.

 

Boiling point

Boiling occurs when molecules within the liquid have the required amount of energy to overcome attractive forces – gas forms inside the liquid and rises to the top due to its density

·        The boiling point is the temperature at which the vapor pressure is equal to the atmospheric pressure – liquid and gas exist in equilibrium

·        At higher elevation, atmospheric pressure decreases and boiling points also decrease

·        In a pressure cooker, atmospheric pressure is increased and boiling point increases

Melting point

            The temperature at which solid and liquid exist in equilibrium – molecules have enough energy to partially overcome the attractive forces locking them into place.

 

Heat transfer and changes of state

·        As a solid absorbs energy its temperature rises until it arrives at the melting point.

·        At this point energy continues to be absorbed, with no change in temperature – energy is being used to break attractive forces.

·        As a liquid absorbs energy its temperature rises until it arrives at the boiling point.

·        At this point energy continues to be absorbed, with no change in temperature – energy is being used to break attractive forces.

·        As a gas absorbs energy its temperature rises.

·        Each substance has a specific heat capacity / molar heat capacity dependent on the state of matter, a molar heat of vaporization, and a molar heat of fusion.

·        See p. 449 and in-class handout

 

How much energy would be required to raise 100 grams of ice from -7°C to 104°C?

 

-7°C (s) à 0°C (s)

 

 

0°C (s) à 0°C (l)

 

 

0°C (l) à 100°C (l)

 

 

100°C (l) à 100°C (g)

 

 

100°C (g) à 104°C (g)

 

 

 

Phase Diagrams

 

Phase diagrams show a plot of P vs. T

·        Lines and curves on the diagram represent the pressure-temperature conditions under which equilibria between states of matter exist

·        Possible equilibria: solid ßà liquid;     

liquidßà gas;                         solid ßà gas

·        The areas between lines/curves are conditions when the sample is a solid, liquid, or gas.

·        The slope of the solid-liquid line describes whether the solid or the liquid is more dense

·        In most cases the solid is more dense – solid particles are packed together more closely than are liquid particles

·        For water, the solid is less dense – an increase in pressure can cause solid ice to melt

·        The network of hydrogen bonding in solid ice is rigid and creates “holes” between H2O particles which decrease the density slightly

·        Solid ice floats in water due to the difference in density

·        The point where all three curves meet is called the triple point – all three phases of matter exist in equilibria under these conditions

·        At pressures below the triple point pressure, the liquid state cannot exist – the substance moves directly from solid to gas as temperature increases (sublimation)

·        Dry ice (solid CO2) sublimes at atmospheric pressure and at room temperature

·        The critical temperature is a temperature above which gas cannot be liquefied, no matter how much pressure is increased

·        The critical pressure is the pressure required to liquefy a gas at the critical temperature

 

Refer to the phase diagrams on page 492 & 497 for the following questions.

·        What state of matter exists for H2O at:

-10°C and 218 atm                     25°C and 350 Torr

·        What state of matter exists for CO2 at:

-80°C and 1 atm                         25°C and 73 atm

·        What transition occurs for CO2 as temperature is increased from -80°C to 25°C at a pressure of 1 atm?

 

Solids can be amorphous (randomly arranged particles) or crystalline (particles arranged in a repeating pattern)

·        Crystalline solids contain repeating units of particles – unit cells

·        Unit cells are described by the length of its edges and the angles between the edges

·        Common arrangements include: cubic, tetragonal, orthorhombic, monoclinic, rhombohedral, hexagonal, and triclinic

·        Atoms, ions, or molecules are found at each of the corners of the cell and are shared by other cells

·        Identical atoms, ions, or molecules may also be found at other locations in the cell – see body-centered cubic and face-centered cubic.

·        Unit cells stack in three dimensions to build a lattice

 

Knowing the dimensions of the cell and the type of substance can allow for a calculation of density (Density = mass/volume)

·        Gold crystallizes as a face-centered cubic cell with a length on each side of 4.10 A.  What is the density of solid gold?

8 ´ 1/8 atoms + 3 ´ ½ atom = 4 atoms/cell

 

4 atoms ´                         1 mol               ´         196.97 g         =

                              6.022 x 1023 atom                1 mol

 

1 A = 1 ´ 10-10 m = 1 ´ 10-8 cm

 

(4.10 A)3   ´ (1 ´ 10-8 cm)3         =

                              1 A3

 

Density = 1.308 ´ 10-21 g/6.89 ´ 10-23 cm3

                              =          19.0 g/cm3

 

 

 

 

 

 

Molecular compounds can also exhibit crystalline structures, with molecules existing at the corners of the unit cell

·        Intermolecular forces such as hydrogen bonding, dipole-dipole forces, or London forces hold the individual molecules in their positions within the crystalline structure

(ex. Solid H2O as shown on p. 460)

 

Covalent solids refers to networks of atoms held in a crystalline structure by covalent bonds

·        Graphite and diamonds are examples of how carbon can form different covalent solids with different patterns (see p. 468).

·        Since the atoms are held in place by covalent bonds rather than intermolecular forces, these compounds have much higher melting points than do molecular solids.