Liquids and Solids Chapter 10

Forces that occur between molecules. 10.1 Intermolecular Forces Forces that occur between molecules. Dipole-Dipole Hydrogen Bonding London Dispersion IMFs are weaker than intramolecular bonds.

10.1 Intermolecular Forces Hydrogen Bonding – occurs between hydrogen and highly electronegative elements (H2O) and can include lone pair electrons (NH3)

10.1 Intermolecular Forces Phase Changes Changes in state are due to changes in the forces between molecules. IMFs still exists, but their ability to be maintained changes during phase changes. The conditions under which a substance changes phase is directly related to IMFs

Less IMF influence Smaller Density Greater Disorder 10.1 Intermolecular Forces Less IMF influence Smaller Density Greater Disorder

Dipole-Dipole Forces Hydrogen Bonds 10.1 Intermolecular Forces Dipole-Dipole Forces Molecules with polar bonds with partial positive and partial negative regions align end to end. Creates greater intermolecular attraction. 1% the strength of covalent bonds. Hydrogen Bonds Special case involving H, O, N, and halogens

Hydrogen bonding is responsible for the crystal structure of ice. 10.1 Intermolecular Forces Ice Hydrogen bonding is responsible for the crystal structure of ice. What is in the large gaps in the crystal structure?

London Dispersion Forces 10.1 Intermolecular Forces London Dispersion Forces Spontaneous dipole induced by random electron motion. Can induce near by molecules to form a dipole. Significant in large atoms (more electrons) Occurs in all molecules Best examples are Br2and I2

Boiling point and freezing point are directly related to IMFs. 10.1 Intermolecular Forces Boiling point and freezing point are directly related to IMFs.

Reality Check Which molecule has a higher boiling point? Ethanol 10.1 Intermolecular Forces Reality Check Which molecule has a higher boiling point? Ethanol Dimethyl Ether

e) All of the above have the same boiling point. 10.1 Intermolecular Forces Reality Check Which of the following would you expect to have the highest boiling point? a) F2 b) Cl2 c) Br2 d) I2 e) All of the above have the same boiling point.

10.2 Liquid State Liquids Low compressibility, lack of rigidity, and high density compared to gases. Surface Tension: resistance of a liquid to increase its surface area. Liquids with strong IMFs tend to have high surface tensions This is why water can be “hard”, why insects can rest on the surface of water, and why water beads on surfaces that are non-polar.

10.2 Liquid State Liquids Capillary action: spontaneous rising of liquid in a narrow tube: Cohesive Forces: intermolecular forces among the molecules of a liquid Adhesive Forces: forces between the molecules and their container. Viscosity: Ability to flow. Liquids with strong IMFs have a high viscosity (low ability to flow) Water forms a concave meniscus because of adhesion forces

Solids Amorphous Solids Crystalline Solids Disorder in the structures 10.3 Introduction to Structures and Types of Solids Solids Amorphous Solids Disorder in the structures Glass Crystalline Solids Ordered Structures Unit Cells

Three cubic unit cells and the corresponding lattices. 10.3 Introduction to Structures and Types of Solids Three cubic unit cells and the corresponding lattices.

λ = wavelength of the X rays d = distance between the atoms 10.3 Introduction to Structures and Types of Solids Bragg Equation, which is used to determine interatomic spacing. n = integer λ = wavelength of the X rays d = distance between the atoms Θ = angle of incidence and reflection

Types of Crystalline Solids 10.3 Introduction to Structures and Types of Solids Types of Crystalline Solids Ionic Solids – ions at the points of the lattice that descries the structure of the solid. Molecular Solids – discrete covalently bonded molecules at each of its lattice points. Atomic Solids – atoms at the lattice points that describe the structure of the solid.

Examples of Three Types of Crystalline Solids 10.3 Introduction to Structures and Types of Solids Examples of Three Types of Crystalline Solids

10.4 Structure and Bonding in Metals Closest Packing Model Assumes that metal atoms are uniform, hard sphere. Spheres are packed in layers. aba packing – the 2nd layer is like 1st, but it is displaced so that each sphere in the 2nd layer occupies a dimple in the 1st layer. The spheres in the 3rd layer occupy the dimples in the 2nd layer so that the spheres in the 3rd layer lie directly over those in the 1st layer

The Closest Packing Arrangement of Uniform Sphere. 10.4 Structure and Bonding in Metals The Closest Packing Arrangement of Uniform Sphere. abc packing: the spheres in the 3rd layer occupy dimples in the 2nd layer so that no spheres in the 3rd layer lie above any layer in the 1st layer. The 4th layer is like the 1st.

Hexagonal Closest Packing 10.4 Structure and Bonding in Metals Hexagonal Closest Packing

10.4 Structure and Bonding in Metals Cubic Closest Packing

The indicated sphere has 12 nearest neighbors 10.4 Structure and Bonding in Metals The indicated sphere has 12 nearest neighbors Each sphere in both ccp and hcp has 12 equivalent nearest neighbors.

Counting parts of atoms – Face Centered Cubic Unit Cell 10.4 Structure and Bonding in Metals Counting parts of atoms – Face Centered Cubic Unit Cell

10.4 Structure and Bonding in Metals Reality Check Determine the number of metal atoms in a unit cell if the packing is: Simple Cubic Cubic Closest packing -1 metal atom -4 metal atoms

Calculate the density of the silver metal. Density = 10.5 g/cm3 10.4 Structure and Bonding in Metals Silver metal crystallizes in a cubic closest packed structure. The face centered cubic unit cell edge is 409 pm. Calculate the density of the silver metal. Density = 10.5 g/cm3 The density is 10.5 g/cm3. There are four silver atoms per unit cell. The volume of the unit cell is d3, which is (409 pm)3, or 6.84 × 107 pm3. This is equivalent to 6.84 × 10-23 cm3. Density = mass/volume = [(4 atoms)(107.9 g/mol)(1 mol/6.022×1023 atoms)] / 6.84 × 10-23 cm3 = 10.5 g/cm3

10.4 Structure and Bonding in Metals The Electron Sea Model A regular array of cations in a “sea” of mobile valence electrons.

10.4 Structure and Bonding in Metals Metal Alloys Substitutional Alloy: some of the host metal atoms are replaced by other metal atoms of similar size. Interstitial Alloy: some of the holes in the closest packed metal structure are occupied by small atoms.

10.5 Carbon and Silicon: Network Atomic Solids

The p Orbitals and Pi-system in Graphite 10.5 Carbon and Silicon: Network Atomic Solids The p Orbitals and Pi-system in Graphite

This is a great section. You really should read it. 10.6 Molecular Solids This is a great section. You really should read it.

Hole size: trigonal < tetrahedral < octahedral 10.7 Ionic Solids Ionic solids can have three types of holes based on the ion arrangement Hole size: trigonal < tetrahedral < octahedral

10.7 Ionic Solids A helpful table

Behavior of a liquid in a closed container 10.8 Vapor Pressure and Changes of State Behavior of a liquid in a closed container Initially At Equilibrium

10.8 Vapor Pressure and Changes of State

Vapor Pressure Pressure of the vapor present at equilibrium 10.8 Vapor Pressure and Changes of State Vapor Pressure Pressure of the vapor present at equilibrium The system is at equilibrium when not net change occurs in the amount of liquid or vapor because the two opposite processes exactly balance each other. Strong IMF = low vapor pressures (WHY?) Vapor pressure increases significantly with temperature.

10.8 Vapor Pressure and Changes of State

Clausius-Clapeyron Equation 10.8 Vapor Pressure and Phase Change Clausius-Clapeyron Equation Pvap = vapor pressure ΔHvap = enthalpy of vaporization R = 8.3145 J/K·mol T = temperature (in kelvin)

Heating Curve for Water 10.8 Vapor Pressure and Changes of State Heating Curve for Water

10.9 Phase Diagrams Graph representing phase changes as a function of temperature and pressure: Triple Point Critical Point Phase Equilibrium Lines

Phase diagram for carbon dioxide 10.9 Phase Diagrams Phase diagram for carbon dioxide

Phase diagram for water 10.9 Phase Diagrams Phase diagram for water

10.9 Phase Changes As intermolecular forces increase, what happens to each of the following? Why? Boiling point Viscosity Surface tension Enthalpy of fusion Freezing point Vapor pressure Heat of vaporization Boiling point increases Viscosity increases Surface tension increases Enthalpy of fusion increases Freezing point increases Vapor pressure decreases Heat of vaporization increases

10.1 Intermolecular Forces

10.1 Intermolecular Forces