Chemistry: The Molecular Nature of Matter, 6E Jespersen/Brady/Hyslop

Chemistry: The Molecular Nature of Matter, 6E Jespersen/Brady/Hyslop
Chapter 12: Intermolecular Attractions and the Properties of Liquids and Solids Chemistry: The Molecular Nature of Matter, 6E Jespersen/Brady/Hyslop

Intermolecular Forces
Important differences between gases, solids, and liquids: Gases Expand to fill their container Liquids Retain volume, but not shape Solids Retain volume and shape

Physical state of molecule depends on Average kinetic energy of particles Recall KE  Tave Intermolecular Forces Energy of Inter-particle attraction Physical properties of gases, liquids and solids determined by How tightly molecules are packed together Strength of attractions between molecules

Intermolecular Attractions
Converting gas  liquid or solid Molecules must get closer together Cool or compress Converting liquid or solid  gas Requires molecules to move farther apart Heat or reduce pressure As T decreases, kinetic energy of molecules decreases At certain T, molecules don’t have enough energy to break away from one another’s attraction

Inter vs. Intra-Molecular Forces
Covalent bonds within molecule Strong Hbond (HCl) = 431 kJ/mol Intermolecular forces Attraction forces between molecules Weak Hvaporization (HCl) = 16 kJ/mol Covalent Bond (strong) Intermolecular attraction (weak)

Electronegativity Review
Electronegativity: Measure of attractive force that one atom in a covalent bond has for electrons of the bond

Bond Dipoles Two atoms with different electronegativity values share electrons unequally Electron density is uneven Higher charge concentration around more electronegative atom Bond dipoles Indicated with delta (δ) notation Indicates partial charge has arisen

Net Dipoles Symmetrical molecules Asymmetrical molecules
Even if they have polar bonds Are non-polar because bond dipoles cancel Asymmetrical molecules Are polar because bond dipoles do not cancel These molecules have permanent, net dipoles Molecular dipoles Cause molecules to interact Decreased distance between molecules increases amount of interaction

When substance melts or boils Intermolecular forces are broken Not covalent bonds Responsible for non-ideal behavior of gases Responsible for existence of condensed states of matter Responsible for bulk properties of matter Boiling points and melting points Reflect strength of intermolecular forces

Three Important Types of Intermolecular Forces
London dispersion forces Dipole-dipole forces Hydrogen bonds Ion-dipole forces Ion-induced dipole forces

London Forces When atoms near one another, their valence electrons interact Repulsion causes electron clouds in each to distort and polarize Instantaneous dipoles result from this distortion Effect enhanced with increased volume of electron cloud size Effect diminished by increased distance between particles and compact arrangement of atoms

London Forces Instantaneous dipole-induced dipole attractions
Dispersion forces Operate between all molecules Neutral or net charged Nonpolar or polar

London Dispersion Forces
Ease with which dipole moments can be induced and thus London Forces depend on Polarizability of electron cloud Points of attraction Number atoms Molecular shape (compact or elongated)

Polarizability Ease with which the electron cloud can be distorted
Larger molecules often more polarizable Larger number of less tightly held electrons Magnitude of resulting partial charge is larger Larger electron cloud

Table 12.1 Boiling Points of Halogens and Noble Gases
Larger molecules have stronger London forces and thus higher boiling points.

Number of Atoms in Molecule
London forces depend on number atoms in molecule Boiling point of hydrocarbons demonstrates this trend Formula BP at 1 atm, C CH4 –161.5 C5H12 36.1 C2H6 –88.6 C6H14 68.7 C3H8 –42.1 : C4H10 –0.5 C22H46 327

How Intermolecular Forces Determine Physical Properties
Hexane, C6H14 BP 68.7 °C More sites (marked with *) along its chain where attraction to other molecules can occur Propane, C3H8 BP –42.1 °C

Molecular Shape Increased surface area available for contact = increased London forces London dispersion forces between spherical molecules are lower than chain-like molecules More compact molecules Hydrogen atoms not as free to interact with hydrogen atoms on other molecules Less compact molecules Hydrogen atoms have more chance to interact with hydrogen atoms on other molecules

Physical Origin of Shape Effect
Small area for interaction Larger area for interaction More compact – lower BP Less compact – higher BP

Dipole-Dipole Attractions
Occur only between polar molecules Possess dipole moments Molecules need to be close together Polar molecules tend to align their partial charges Positive to negative As dipole moment increases, intermolecular force increases      

Tumbling molecules Mixture of attractive and repulsive dipole-dipole forces Attractions (- -) are maintained longer than repulsions(- -) Get net attraction ~1–4% of covalent bond

Interactions between net dipoles in polar molecules About 1–4% as strong as a covalent bond Decrease as molecular distance increases Dipole-dipole forces increase with increasing polarity

Hydrogen Bonds Special type of dipole-dipole Interaction
Very strong dipole-dipole attraction ~10% of a covalent bond Occurs between H and highly electronegative atom (O, N, or F) H—F, H—O, and H—N bonds very polar Electrons are drawn away from H, so high partial charges H only has one electron, so +H presents almost bare proton –X almost full –1 charge Element’s small size, means high charge density Positive end of one can get very close to negative end of another

Examples of Hydrogen Bonding

Hydrogen Bonding in Water
Responsible for expansion of water as it freezes Hydrogen bonding produces strong attractions in liquid Hydrogen bonding (dotted lines) between water molecules in ice form tetrahedral configuration

Your Turn! List all intermolecular forces for CH3CH2OH. A. Hydrogen-bonds B. Hydrogen-bonds, dipole-dipole attractions, London dispersion forces C. Dipole-dipole attractions D. London dispersion forces E. London dispersion forces, dipole-dipole attractions

Your Turn! In the liquid state, which species has the strongest intermolecular forces, CH4, Cl2, O2 or HF? A. CH4 B. Cl2 C. O2 D. HF

Ion-Dipole Attractions
Attractions between ion and charged end of polar molecules Attractions can be quite strong as ions have full charges (a) Negative ends of water dipoles surround cation (b) Positive ends of water dipoles surround anion

Ex. Ion-Dipole Attractions: AlCl3·6H2O
Attractions between ion and polar molecules Positive charge of Al3+ ion attracts partial negative charges – on O of water molecules Ion-dipole attractions hold water molecules to metal ion in hydrate Water molecules are found at vertices of octahedron around aluminum ion

Ion-Induced Dipole Attractions
Attractions between ion and dipole it induces on neighboring molecules Depends on Ion charge and Polarizability of its neighbor Attractions can be quite strong as ion charge is constant, unlike instantaneous dipoles of ordinary London forces E.g., I– and Benzene

Summary of Intermolecular Attractions
Dipole-dipole Occur between neutral molecules with permanent dipoles About 1–4% of covalent bond Mid range in terms of intermolecular forces Hydrogen bonding Special type of dipole-dipole interaction Occur when molecules contain N—H, H—F and O—H bonds About 10% of a covalent bond

Summary of Intermolecular Attractions
London dispersion Present in all substances Weakest intermolecular force Weak, but can add up to large net attractions Ion-dipole Occur when ions interact with polar molecules Strongest intermolecular attraction Ion-induced dipole Occur when ion induces dipole on neighboring particle Depend on ion charge and polarizability of its neighbor

Using Intermolecular Forces
Often can predict physical properties (like BP, MP and many others) by comparing strengths of intermolecular attractions Ion-Dipole Hydrogen Bonding Dipole-Dipole London Forces Larger, longer, and therefore heavier molecules often have stronger intermolecular forces Smaller, more compact, lighter molecules have generally weaker intermolecular forces Strongest Weakest

Physical Properties that Depend on How Tightly Molecules Pack
Compressibility Measure of ability of substance to be forced into smaller volume Determined by strength of intermolecular forces Gases highly compressible Molecules far apart Weak intermolecular forces Solids and liquids nearly incompressible Molecules very close together Stronger intermolecular forces

Intermolecular Forces Determine Strength of Many Physical Properties
Retention of volume and shape Solids retain both volume and shape Strongest intermolecular attractions Molecules closest Liquids retain volume, but not shape Attractions intermediate Gases, expand to fill their containers Weakest intermolecular attractions Molecules farthest apart

Intermolecular Forces and Temperature
Decrease with increasing temperature Increasing kinetic energy overcomes attractive forces If allowed to expand, increasing temperature increases distance between gas particles and decreases attractive forces

Diffusion Movement that spreads one gas though another gas to occupy space uniformly Spontaneous intermingling of molecules of one gas with molecules of another gas Occurs more rapidly in gases than in liquids Hardly at all in solids

Diffusion In Gases In Liquids In Solids
Molecules travel long distances between collisions Diffusion rapid In Liquids Molecules closer Encounter more collisions Takes a long time to move from place to place In Solids Diffusion close to zero at room temperature Will increase at high temperature Diffusion

Surface Tension Why does H2O bead up on a freshly waxed car instead of forming a layer? Inside body of liquid Intermolecular forces are the same in all directions Molecules at surface Potential energy increases when removing neighbors Molecules move together to reduce surface area and potential energy

Surface Tension Causes a liquid to take the shape (a sphere) that minimizes its surface area Molecules at surface have higher potential energy than those in bulk of liquid and move to reduce the potential energy Wax = nonpolar H2O = polar Water beads in order to reduce potential energy by reducing surface area

Surface Tension Liquids containing molecules with strong intermolecular forces have high surface tension Allows us to fill glass above rim Gives surface rounded appearance Surface acts as “skin” that lets water pile up Surface resists expansion and pushes back Surface tension increases as intermolecular forces increase Surface tension decreases as temperature increases

Wetting Ability of liquid to spread across surface to form thin film
Greater similarity in attractive forces between liquid and surface, yields greater wetting effect Occurs only if intermolecular attractive force between surface and liquid about as strong as within liquid itself

Wetting Ex. H2O wets clean glass surface as it forms
H–bonds to SiO2 surface Does not wet greasy glass, because grease is nonpolar and water is very polar Only London forces Forms beads instead Surfactants Added to detergents to lower surface tension of H2O Now water can spread out on greasy glass

Surfactants (Detergents)
Substances that have both polar and non-polar characteristics Long chain hydrocarbons with polar tail Nonpolar end dissolves in nonpolar grease Polar end dissolves in polar H2O Thus increasing solubility of grease in water

Viscosity Resistance to flow
Measure of fluid’s resistance to flow or changing form Related to intermolecular attractive forces Also called internal friction Depends on intermolecular attractions

Viscosity Viscosity decreases when temperature increases
Most people associate liquids with viscosity Syrup more viscous than water Gases have viscosity Respond almost instantly to form-changing forces Solids, such as rocks and glass have viscosity Normally respond very slowly to forces acting to change their shape

Effect of Intermolecular Forces on Viscosity
Ethylene glycol Polar molecule Hydrogen-bonding Dipole-dipole and London forces Acetone Polar molecule Dipole-dipole and London forces Which is more viscous?

Your Turn! For each pair given, which is has more viscosity? CH3CH2CH2CH2OH, CH3CH2CH2CHO C6H14, C12H26 NH3(l ), PH3(l ) A. CH3CH2CH2CH2OH C6H14 NH3(l ) B. CH3CH2CH2CH2OH C12H26 NH3(l ) C. CH3CH2CH2CHO C6H14 PH3(l ) D. CH3CH2CH2CHO C12H26 NH3(l ) E. CH3CH2CH2CH2OH C12H26 PH3(l )

Solubility “Like dissolves like” Compare relative polarity Surfactants
To dissolve polar substance, use polar solvent To dissolve nonpolar substance, use nonpolar solvent Compare relative polarity Similar polarity means greater ability to dissolve in each other Differing polarity means that they don’t dissolve, they are insoluble Surfactants Both polar and non-polar characteristics Used to increase solubility

Your Turn! Which of the following are not expected to be soluble in water? HF CH4 CH3OH All are soluble CH4, methane, is the only nonpolar substance so only has London forces. HF and CH3OH are both polar molecules and have hydrogen bonds.

Phase Changes Changes of physical state As temperature changes
Deal with motion of molecules As temperature changes Matter will undergo phase changes Liquid  Gas Evaporation, vaporization As heat is added, H2O, forms steam or water vapor Requires energy or source of heat

Phase Changes Solid  Gas Gas  Liquid Sublimation
Ice cubes in freezer, leave in long enough disappear Endothermic Gas  Liquid Condensation Dew is H2O vapor condensing onto cooler ground Exothermic Often limits lower night time temperature

Rate of Evaporation Depends on Temperature Surface area
Strength of intermolecular attractions Molecules that escape from liquid have larger than minimum escape KE When they leave Average KE of remaining molecules is less and so T lower

Effect of Temperature on Evaporation Rate
For given liquid Rate of evaporation per unit surface area increases as T increases Why? At higher T, total fraction of molecules with KE large enough to escape is larger Result: rate of evaporation is larger

Kinetic Energy Distribution in Two Different Liquids
A B Smaller intermolecular forces Lower KE required to escape liquid A evaporates faster Larger intermolecular forces Higher KE required to escape liquid B evaporates slower

Changes Of State Involve Equilibria
Fraction of molecules in condensed state is higher when intermolecular attractions are higher Intermolecular attractions must be overcome to separate the particles, while separated particles are simultaneously attracted to one another

Before System Reaches Equilibrium
Liquid is placed in empty, closed, container Begins to evaporate Once in gas phase Molecules can condense by Striking surface of liquid and giving up some kinetic energy

System At Equilibrium Rate of evaporation = rate of condensation
Occurs in closed systems where molecules cannot escape

Similar Equilibria Reached in Melting
Melting Point (mp) Solid begins to change into liquid as heat added Dynamic equilibria exists between solid and liquid states Melting (red arrows) and freezing (black arrows) occur at same rate As long as no heat added or removed from equilibrium mixture

Equilibria Reached in Sublimation
At equilibrium Molecules sublime from solid at same rate as molecules condense from vapor

Phase Changes Gas Energy of System Liquid Solid Vaporization
Condensation Sublimation Deposition Energy of System Liquid Melting or Fusion Freezing Solid  Exothermic, releases heat  Endothermic, absorbs heat

Energy Changes Accompanying Phase Changes
All phase changes are possible under the right conditions Following sequence is endothermic heat solid  melt  heat liquid  boil  heat gas Following sequence is exothermic cool gas  condense  cool liquid  freeze  cool solid

Enthalpy Of Phase Changes
Endothermic Phase Changes Must add heat Energy entering system (+) Sublimation: Hsub > 0 Vaporization: Hvap > 0 Melting or Fusion: Hfus > 0 Exothermic Phase Changes Must give off heat Energy leaving system (–) Deposition: H < 0 = –Hsub Condensation: H < 0 = –Hvap Freezing: H < 0 = –Hfus

Phase Changes As T changes, matter undergoes phase changes
Transformation from one phase to another Liquid-Vapor Equilibrium Molecules in liquid Not in rigid lattice In constant motion Denser than gas, so more collisions Some have enough kinetic energy to escape, some don’t

Liquid-Vapor Equilibrium
At any given T, Average kinetic energy of molecules is constant But particles have a distribution of kinetic energies Certain number of molecules have enough KE to escape surface Fraction of molecules Kinetic Energy As T increases, average KE increases and number molecules with enough KE to escape increases

Vapor Pressure Equilibrium Vapor Pressure
Pressure molecules exert when they evaporate or escape into gas (vapor) phase Pressure of gas when liquid or solid is at equilibrium with its gas phase Increasing temperature increases vapor pressure because vaporization is endothermic liquid + heat of vaporization ↔ gas Equilibrium Vapor Pressure VP once dynamic equilibrium reached Usually referred to as simply vapor pressure

Measuring Vapor Pressure
To measure pressures inside vessels, a manometer is used.

Vapor Pressure Diagram
Variation of vapor pressure with T Ether Volatile High vapor pressure near RT Propylene glycol Non-volatile Low vapor pressure near RT RT = 25 C

Effect of Volume on VP Initial V Increase V More liquid evaporates
Liquid – vapor equilibrium exists Increase V Pressure decreases Rate of condensation decreases More liquid evaporates New equilibrium established

Measuring Hvap Clausius-Clapeyron equation
Measure pressure at various temperatures, then plot Two point form of Clausius-Clapeyron equation Measure pressure at two temperatures and solve equation

Learning Check The vapor pressure of diethyl ether is 401 mm Hg at 18 °C, and its molar heat of vaporization is 26 kJ/mol. Calculate its vapor pressure at 32 °C. T1 = = K T2 = = K

Your Turn! Determine the enthalpy of vaporization, in kJ/mol, for benzene, using the following vapor pressure data. T = 60.6 °C; P = 400 torr T = 80.1 °C; P = 760 torr A kJ/mol B kJ/mol C. –32.4 kJ/mol D kJ/mol E. –14.0 kJ/mol

Your Turn! - Solution

Do Solids Have Vapor Pressures?
Yes At given temperature Some solid particles have enough KE to escape into vapor phase When vapor particles collide with surface They can be captured Equilibrium vapor pressure of solid Pressure of vapor in equilibrium with solid

Boiling Point (bp) Normal Boiling Point
T at which vapor pressure of liquid = atmospheric pressure. Bp increases as strength of intermolecular forces increase Normal Boiling Point T at which vapor pressure of liquid = 1 atm

Effects of Hydrogen Bonding
Boiling points of hydrogen compounds of elements of Groups 4A, 5A, 6A, and 7A. Boiling points of molecules with hydrogen bonding are much higher than expected

Your Turn! Which of the following will affect the boiling point of a substance? Polarizability Intermolecular attractions The external pressure on the material All of these None of these Molecular mass, strength of intermolecular forces and the external pressure on the material all effect the boiling point of a substance.

Heating Curve Heat added at constant rate Horizontal lines
Phase changes Melting point Boiling point Diagonal lines Heating of solid, liquid or gas

Cooling Curve Heat removed at constant rate Horizontal lines
Phase changes Melting point Boiling point Diagonal lines Cooling of solid, liquid or gas Supercooling Temperature of liquid dips below its freezing point

Your Turn! How much heat, in J, is required to convert g of ice at °C to water at °C? Specific heat (J/g K): ice, 2.108, water, Enthalpy of fusion = kJ/mol A J B J C J D J E. 62,400 J

Energies of Phase Changes
Expressed per mole Molar heat of fusion (Hfus) Heat absorbed by one mole of solid when it melts to give liquid at constantT and P Molar heat of vaporization (Hvap ) Heat absorbed when one mole of liquid is changed to one mole of vapor at constant T and P Molar heat of sublimation (Hsub ) Heat absorbed by one mole of solid when it sublimes to give one mole of vapor at constant T and P All of these quantities tend to increase with increasing intermolecular forces

Le Chatelier’s Principle
Equilibria are often disturbed or upset When dynamic equilibrium of system is upset by a disturbance System responds in direction that tends to counteract disturbance and, if possible, restore equilibrium Position of equilibrium Used to refer to relative amounts of substance on each side of double (equilibrium) arrows

Liquid Vapor Equilibrium
Liquid Heat  Vapor Increasing T Increases amount of vapor Decreases amount of liquid Equilibrium has shifted Shifted to the right More vapor is produced at expense of liquid Temperature-pressure relationships can be represented using a phase diagram

4/21/2017 Phase Diagrams Show the effects of both pressure and temperature on phase changes Boundaries between phases indicate equilibrium Triple point: The temperature and pressure at which s, l, and g are all at equilibrium Critical point: The temperature and pressure at which a gas can no longer be condensed TC = temperature at critical point PC = pressure at critical point Chem FAQ: How can I determine what phase will be present at a given pressure and temperature using a phase diagram? Safety Dog Productions

Phase Diagram X axis – temperature Y axis – pressure
As P increases (T constant), solid most likely More compact As T increases (P constant), gas most likely Higher energy Each point = T and P B = E = F = E F 0.01 °C, 4.58 torr 100 °C, 760 torr –10 °C, 2.15 torr

Phase Diagram of Water AB = vapor pressure curve for ice
BD = vapor pressure curve for liquid water BC = melting point line B = triple point: T and P where all three phases are in equilibrium D = critical point T and P above which liquid does not exist

Case Study: An Ice Necklace
A cube of ice may be suspended on a string simply by pressing the string into the ice cube. As the string is pressed onto the surface, it becomes embedded into the ice. Why does this happen?

Phase Diagram – CO2 Now line between solid and liquid slants to right
More typical Where is triple point? Where is critical point?

Supercritical Fluid Substance with temperature above its critical temperature (TC) and density near its liquid density Have unique properties that make them excellent solvents Values of TC tend to increase with increased intermolecular attractions between particles

Your Turn! At 89 °C and 760 mmHg, what physical state is present?
Solid Liquid Gas Supercritical fluid Not enough information is given At 89 °C and 760 mmHg we are in the liquid range—above the line BD for gases.

Types of Solids Crystalline Solids Amorphous Solids
Solids with highly regular arrangements of components Amorphous Solids Solids with considerable disorder in their structures

Crystalline Solids Unit Cell Smallest segment that repeats regularly
Smallest repeating unit of lattice Two- dimensional unit cells

Crystal Structures Have Regular Patterns
Lattice Many repeats of unit cell Regular, highly symmetrical system Three (3) dimensional system of points designating positions of components Atoms Ions Molecules

Three Types Of 3-D Unit Cells
Simple cubic Has one host atom at each corner Edge length a = 2r Where r is radius of atom or ion Body-centered cubic (BCC) Has one atom at each corner and one in center Edge length Face-centered cubic (FCC) Has one atom centered in each face, and one at each corner

Close Packing of Spheres
1st layer 2nd layer Most efficient arrangement of spheres in two dimensions Each sphere has 6 nearest neighbors Second layer with atoms in holes on the first layer

Two Ways to Put on Third Layer
Cubic lattice: 3-dimensional arrays 2. Above holes in first layer Remaining holes not covered by second layer 1. Directly above spheres in first layer

3-D Simple Cubic Lattice
Unit Cell Space filling model Portion of lattice—open view

Other Cubic Lattices Face Centered Cubic Body Centered Cubic
Fig 12.37, 38, 39

Ionic Solids Lattices of alternating charges
Want cations next to anions Maximizes electrostatic attractive forces Minimizes electrostatic repulsions Based on one of three basic lattices: Simple cubic Face centered cubic Body centered cubic

Common Ionic Solids Rock salt or NaCl
Face centered cubic lattice of Cl– ions (green) Na+ ions (blue) in all octahedral holes

Other Common Ionic Solids
Cesium Chloride, CsCl Zinc Sulfide, ZnS Calcium Fluoride, CaF2

Spaces In Ionic Solids Are Filled With Counter Ions
In NaCl Cl– ions form face-centered cubic unit cell Smaller Na+ ions fill spaces between Cl–ions Count atoms in unit cell Have 6 of each or 1:1 Na+:Cl– ratio

Counting Atoms per Unit Cell
Four types of sites in unit cell Central or body position – atom is completely contained in one unit cell Face site – atom on face shared by two unit cells Edge site – atom on edge shared by four unit cells Corner site – atom on corner shared by eight unit cells Site Counts as Shared by X unit cells Body 1 Face 1/2 2 Edge 1/4 4 Corner 1/8 8

Example: NaCl Site # of Na+ # of Cl– Body 1 Face Edge Corner Total 4
Face Edge Corner Total 4 Center

Learning Check: Determine the number of each type of ion in the unit cell. 1:1 CsCl 4:4 ZnS 4:8 CaF2

Some Factors Affecting Crystalline Structure
Size of atoms or ions involved Stoichiometry of salt Materials involved Some substances do not form crystalline solids

Amorphous Solids (Glass)
Have little order, thus referred to as “super cooled liquids” Edges are not clean, but ragged due to the lack of order

X-Ray Crystallography
X rays are passed through crystalline solid Some x rays are absorbed, most re-emitted in all directions Some emissions by atoms are in phase, others out of phase Emission is recorded on film

X-ray Diffraction Experimental Setup Diffraction Pattern

Interpreting Diffraction Data
As x rays hit atoms in lattice they are deflected Angles of deflections related to lattice spacing So we can estimate atomic and ionic radii from distance data

Interpreting Diffraction Data
Bragg Equation nλ=2d sinθ n = integer (1, 2, …)  = wavelength of X rays d = interplane spacing in crystal  = angle of incidence and angle of reflectance of X rays to various crystal planes

Example: Diffraction Data
The diffraction pattern of copper metal was measured with X-ray radiation of wavelength of pm. The first order (n = 1) Bragg diffraction peak was found at an angle θ of 50.5°. Calculate the spacing between the diffracting planes in the copper metal. n = 2d sin  1(131.5 pm) = 2 × d × sin(50.5) d = 283 pm

Example: Using Diffraction Data
X-ray diffraction measurements reveal that copper crystallizes with a face-centered cubic lattice in which the unit cell length is 362 pm. What is the radius of a copper atom expressed in picometers? This is basically a geometry problem.

Ex. Using Diffraction Data (cont.)
Pythagorean theorem: a2 + b2 = c2 Where a = b = 362 pm sides and c = diagonal 2a2 = c and diagonal = 4  rCu = 512 pm rCu = 128 pm

Learning Check Silver packs together in a faced center cubic fashion. The interplanar distance, d, corresponds to the length of a side of the unit cell, and is 407 pm. What is the radius of a silver atom? a r = 53.6 pm

Ionic Crystals (e.g. NaCl, NaNO3)
Have cations and anions at lattice sites Are relatively hard Have high melting points Are brittle Have strong attractive forces between ions Do not conduct electricity in their solid states Conduct electricity well when molten

Sample Homework Problem
Potassium chloride crystallizes with the rock salt structure. When bathed in X rays, the layers of atoms corresponding to the surfaces of the unit cell produce a diffracted beam of X rays (λ=154 pm) at an angle of 6.97°. From this, calculate the density of potassium chloride in g/cm3.

Your Turn! Yitterbium crystallizes with a face centered cubic lattice. The atomic radius of yitterbium is 175 pm. Determine the unit cell length. A. 495 pm B. 700 pm C. 350 pm D. 990 pm E. 247 pm

Your Turn! - Solution

Covalent Crystals Lattice positions occupied by atoms that are covalently bonded to other atoms at neighboring lattice sites Also called network solids Interlocking network of covalent bonds extending all directions Covalent crystals tend to Be very hard Have very high melting points Have strong attractions between covalently bonded atoms

Ex. Covalent (Network) Solid
Diamond (all C) Shown SiO2 silicon oxide Alternating Si and O Basis of glass and quartz Silicon carbide (SiC)

Metallic Crystals Simplest models
Lattice positions of metallic crystal occupied by positive ions Cations surrounded by “cloud” of electrons Formed by valence electrons Extends throughout entire solid

Metallic Crystals Conduct heat and electricity
By their movement, electrons transmit kinetic energy rapidly through solid Have the luster characteristically associated with metals When light shines on metal Loosely held electrons vibrate easily Re-emit light with essentially same frequency and intensity

Learning Check: Classify the following in terms of most likely type of solid. Substance ionic molecular covalent metallic X: Pulverizes when struck; non-conductive of heat and electricity Y: White crystalline solid that conducts electrical current when molten or dissolved Z: Shiny, conductive, malleable with high melting temperature   

Your Turn! Molecular crystals can contain all of the listed attraction forces except: Dipole-dipole attractions Electrostatic forces London forces Hydrogen bonding

Chemistry: The Molecular Nature of Matter, 6E Jespersen/Brady/Hyslop

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