Learning objective 2.16: The student is able to explain the properties (phase, vapor pressure, viscosity, etc.) of small and large molecular compounds.

Learning objective 2.16: The student is able to explain the properties (phase, vapor pressure, viscosity, etc.) of small and large molecular compounds in terms of the strengths and types of intermolecular forces.

What to do: Distinguish between intermolecular and intramolecular attractions Put a list of compounds in order of increasing melting point, boiling point, and vapor pressure Distiguish between types of solids Identify properties of solids Use the Clausius-Clapeyron equation to relate temperature to vapor pressure of a substance Determine energy gained or lost when state of matter changes.

Solid, Liquid, or Gas What are three factors determine whether a substance is a solid, a liquid, or a gas: The attractive intermolecular forces between particles that tend to draw the particles together. Temperature: The kinetic energies of the particles (atoms, molecules, or ions) that make up a substance. Kinetic energy tends to keep the particles moving apart. Pressure: pressure is increased or decreased as the volume of a closed container changes

Types of Attractive Forces
There are several types of attractive intermolecular forces: Ionic Ion-dipole forces Dipole-dipole forces Hydrogen bonding Induced-dipole forces Ion-induced Dipole-induced London dispersion forces All of the intermolecular forces that hold a liquid together are called cohesive forces.

Ionic Bonds Electrons are transferred
Electronegativity differences are generally greater than 1.7 The formation of ionic bonds is always exothermic!

Ion-Dipole Forces An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a neutral molecule that has a dipole. Most commonly found in solutions. Especially important for solutions of ionic compounds in polar liquids. Ion-dipole attractions become stronger as either the charge on the ion increases, or as the magnitude of the dipole of the polar molecule increases.

Dipole-Dipole Forces Dipole-dipole forces are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. They are much weaker than ionic or covalent bonds and have a significant effect only when the molecules involved are close together (touching or almost touching).

Hydrogen Bonding Bonding between hydrogen and more electronegative neighboring atoms such as fluorine, oxygen and nitrogen Hydrogen bonding between ammonia and water

Hydrogen Bonding in DNA
Thymine hydrogen bonds to Adenine T A

Hydrogen Bonding in DNA
Cytosine hydrogen bonds to Guanine C G

Ion-Induced Dipole Forces
Induced dipole forces result when an ion or a dipole induces a dipole in an atom or a molecule with no dipole. These are weak forces. Ion-Induced Dipole Forces An ion-induced dipole attraction is a weak attraction that results when the approach of an ion induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.

Dipole-Induced Dipole Forces
A dipole-induced dipole attraction is a weak attraction that results when a polar molecule induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.

London Dispersion Forces
Instantaneous dipole that occurs accidentally in a given atom induces a similar dipole in a neighboring atom. Significant in large atoms/molecules. Occurs in all molecules, including nonpolar ones. Fritz London

London Dispersion Forces

London Forces in Hydrocarbons

Boiling point as a measure of intermolecular attractive forces

Relative Magnitudes of Forces
The types of bonding forces vary in their strength as measured by average bond energy. Ionic bonds Strongest Weakest *Ion-dipole interactions *Hydrogen bonding (12-16 kcal/mol ) Dipole-dipole interactions (2-0.5 kcal/mol) Ion induced dipole interactions Induced Dipole-dipole interactions London forces (less than 1 kcal/mol)

Overview of Particle Forces
Type of Force Relative Strength Present in Examples London Dispersion Force (LD) Weak, increasing with size atoms/molecules H2 (g) Ions or Polar molecules with nonpolar molecules NaCl with I2 weak Induced-dipole force Polar molecules HCl moderate Dipole-dipole force Molecules with hydrogen bonded to N,O, or F Hydrogen bond Strong HF, NH3, H2O Ion-dipole Very strong Mixtures of ionic and polar compounds Na+,Cl-, and H2O Ionic Very very strong Lattice structures Na+and Cl-

CH4 CaCl2 NaHCO3 CH3Cl C12H22O11 NH3 H2S NH4Cl
Identify the predominant intermolecular forces present in the solids of each of the following substances: dipole-dipole ion-dipole hydrogen bonding London dispersion dipole-induced dipole ionic ion-induced dipole CH4 CaCl2 NaHCO3 CH3Cl C12H22O11 NH3 H2S NH4Cl

CH4 CaCl2 NaHCO3 CH3Cl C12H22O11 NH3 H2S NH4Cl (d) (f) (f) (a) (c) (c)
Identify the predominant intermolecular forces present in the solids of each of the following substances: dipole-dipole ion-dipole hydrogen bonding London dispersion dipole-induced dipole ionic ion-induced dipole CH4 CaCl2 NaHCO3 CH3Cl C12H22O11 NH3 H2S NH4Cl (d) (f) (f) (a) (c) (c) (f) (a)

SO2 and CH4 MgCl2 and C3H8 C4H10 and BF3 C12H22O11 and C3H8
Identify the intermolecular forces present between the two substances listed: dipole-dipole ion-dipole hydrogen bonding London dispersion dipole-induced dipole ionic ion-induced dipole SO2 and CH4 MgCl2 and C3H8 C4H10 and BF3 C12H22O11 and C3H8 NBr3 and H2S H2O and C12H22O11 NBr3 and H2O CO2 and CCl4

SO2 and CH4 MgCl2 and C3H8 C4H10 and BF3 NBr3 and H2S
Identify the strongest intermolecular forces present between the two substances listed: dipole-dipole ion-dipole hydrogen bonding London dispersion dipole-induced dipole ionic ion-induced dipole SO2 and CH4 MgCl2 and C3H8 C4H10 and BF3 C12H22O11 and C3H8 NBr3 and H2S H2O and C12H22O11 NBr3 and H2O CO2 and CCl4 (e) (g) (d) (e) (a) (c) (c) (d)

Intermolecular Quiz

CO MgCl2 CO2 CF2Cl2 C6H12O6 NF3 H2O2 NH4Br
Identify the predominant intermolecular forces present in the solids of each of the following substances: dipole-dipole ion-dipole hydrogen bonding London dispersion dipole-induced dipole ionic ion-induced dipole CO MgCl2 CO2 CF2Cl2 C6H12O6 NF3 H2O2 NH4Br

SO2 and CF4 CO2 and CH4 CaCl2 and C2H6 NF3 and H2S H2O and C6H12O6
Identify the intermolecular forces present between the two substances listed: dipole-dipole ion-dipole hydrogen bonding London dispersion dipole-induced dipole ionic ion-induced dipole SO2 and CF4 CO2 and CH4 CaCl2 and C2H6 NF3 and H2S H2O and C6H12O6 NF3 and H2O

a. SO2 b. CF4 a. CaCl2 b. C2H6 a. C2H6 b. C6H12O6 a. H2O b. NaCl c. HF
Identify the substance that would have the highest boiling point for each pair: a. SO b. CF4 a. CaCl b. C2H6 a. C2H6 b. C6H12O6 a. H2O b. NaCl c. HF a. Cl b. Br c. I2 a. HF b. HCl c. HBr

Trade and grade. 1 pt for each.

CO (a) MgCl2 (f) CO2 (d) CF2Cl2 (a) C6H12O6 (c) NF3 (a) H2O2 (c)
Identify the predominant intermolecular forces present in the solids of each of the following substances: dipole-dipole ion-dipole hydrogen bonding London dispersion dipole-induced dipole ionic ion-induced dipole CO (a) MgCl2 (f) CO2 (d) CF2Cl2 (a) C6H12O6 (c) NF3 (a) H2O2 (c) NH4Br (f)

SO2 and CF4 (e) CO2 and CH4 (d) CaCl2 and C2H6 (g) NF3 and H2S (a)
Identify the intermolecular forces present between the two substances listed: dipole-dipole ion-dipole hydrogen bonding London dispersion dipole-induced dipole ionic ion-induced dipole SO2 and CF4 (e) CO2 and CH4 (d) CaCl2 and C2H6 (g) NF3 and H2S (a) H2O and C6H12O6 (c) NF3 and H2O (c)

a. SO2 b. CF4 (a) a. CaCl2 b. C2H6 (a) a. C2H6 b. C6H12O6 (b)
Identify the substance that would have the highest boiling point for each pair: a. SO b. CF4 (a) a. CaCl b. C2H6 (a) a. C2H6 b. C6H12O6 (b) a. H2O b. NaCl c. HF (b) Cl b. Br c. I2 (c) a. HF b. HCl c. HBr (a)

Chromatography Concept
Stationary Phase (the paper) Mobile Phase (solvent traveling up the paper)

What Is a Liquid? A liquid is a state of matter in which a sample of matter: is made up of very small particles (atoms, molecules, and/or ions). flows and can change its shape. is not easily compressible and maintains a relatively fixed volume. The particles that make up a liquid: are close together with no regular arrangement, vibrate, move about, and slide past each other. This bottle contains both liquid bromine [Br2(l), the darker phase at the bottom of the bottle] and gaseous bromine [Br2(g), the lighter phase above the liquid]. The circles show microscopic views of both liquid bromine and gaseous bromine.

More Properties of a Liquid
Surface Tension: The resistance to an increase in its surface area (polar molecules, liquid metals). Capillary Action: Spontaneous rising of a liquid in a narrow tube.

Even More Properties of a Liquid
Viscosity: Resistance to flow Viscosity is the resistance to flow. Viscose liquids slowly flow such as syrup, while non viscose liquids flow rapidly, such as water. High viscosity is an indication of strong intermolecular forces

Evaporation Evaporation is the change of a liquid to a gas. Microscopic view of a liquid. Microscopic view after evaporation. When a liquid is heated sufficiently or when the pressure on the liquid is decreased sufficiently, the forces of attraction between molecules do not prevent them from moving apart, and the liquid evaporates to a gas. Example: The sweat on the outside of a cold glass evaporates when the glass warms. Example: Gaseous carbon dioxide is produced when the valve on a CO2 fire extinguisher is opened and the pressure is reduced.

Condensation Condensation is the change from a vapor to a condensed state (solid or liquid). When a gas is cooled sufficiently or, in many cases, when the pressure on the gas is increased sufficiently, the forces of attraction between molecules prevent them from moving apart, and the gas condenses to either a liquid or a solid. Example: Water vapor condenses and forms liquid water (sweat) on the outside of a cold glass or can. Example: Liquid carbon dioxide forms at the high pressure inside a CO2 fire extinguisher. Microscopic view of a gas. Microscopic view after condensation.

Solids Two major arrangements:
Crystalline- have a regular arrangement of components in their structure. Amorphous- those with much disorder in their structure.

Solids There are many amorphous solids, like glass.
We tend to focus on crystalline solids. Four Categories:

Substitution alloys Interstitial alloys
Atoms of the alloying agent replace atoms of the main metal. In most substitution alloys, the constituent elements are quite near one another in the periodic table, and have atoms of roughly similar size. Brass, for example, is a substitution alloy based on copper in which atoms of zinc replace 10–35 percent of the atoms that would normally be in copper. Interstitial alloys Alloys can also form if the alloying agent or agents have atoms that are very much smaller than those of the main metal. The agent atoms slip in between the main metal atoms (in the gaps or "interstices"). Steel is an example of an interstitial alloy in which a relatively small number of carbon atoms slip in the gaps between the huge atoms in a crystalline lattice of iron.

Covalent Network Solids

Covalent Network Solids Carbon- A Special Atomic Solid
There are three types of solid carbon. Coal - amorphous. Diamond- hardest natural substance on earth, insulates both heat and electricity. Graphite- slippery, conducts electricity.

Diamond- each Carbon is sp3 hybridized, connected to four other carbons.
Carbon atoms are locked into tetrahedral shape. Strong s bonds give the huge molecule its hardness.

Graphite is different. Each carbon is connected to three other carbons and sp2 hybridized. The molecule is flat with 120º angles in fused 6 member rings. The p bonds extend above and below the plane.

Molecular solids. Different molecules or atoms have different forces between them. These forces depend on the size of the molecule or atom. They also depend on the strength and nature of dipole moments.

Those without dipoles. Most are gases at 25ºC.
The only forces are London Dispersion Forces. These depend on size of atom. Large molecules (such as I2 ) can be solids even without dipoles.

Ionic Solids The atoms are actually held together by opposite charges.
Huge melting and boiling points. Atoms are locked in lattice so hard and brittle. Every electron is accounted for so they are poor conductors-good insulators.

Put it to practice with problem 81

Vapor Pressure The vapor pressure of a liquid is the equilibrium pressure of a vapor above its liquid (or solid) The vapor pressure of a liquid varies with its temperature, as the following graph shows for water. The line on the graph shows the boiling temperature for water. As the temperature of a liquid or solid increases its vapor pressure also increases. Conversely, vapor pressure decreases as the temperature decreases.

Factors That Affect Vapor Pressure
Types of Molecules: the types of molecules that make up a solid or liquid determine its vapor pressure. If the intermolecular forces between molecules are: relatively strong, the vapor pressure will be relatively low. relatively weak, the vapor pressure will be relatively high. substance vapor pressure at 25oC diethyl ether 0.7 atm bromine 0.3 atm ethyl alcohol 0.08 atm water 0.03 atm Surface Area: the surface area of the solid or liquid in contact with the gas has no effect on the vapor pressure.

Temperature Dependence of Vapor Pressures
The vapor pressure above the liquid varies exponentially with changes in the temperature. The Clausius-Clapeyron equation shows how the vapor pressure and temperature are related. It can be written as:

Clausius – Clapeyron Equation
A straight line plot results when ln P vs. 1/T is plotted and has a slope of Hvap/R. Clausius – Clapeyron equation is true for any two pairs of points. Write the equation for each and combine to get:

Using the Clausius – Clapeyron Equation
Boiling point - the temperature at which the vapor pressure of a liquid is equal to the pressure of the external atmosphere. Normal boiling point - the temperature at which the vapor pressure of a liquid is equal to atmospheric pressure (1 atm). E.g. Determine normal boiling point of chloroform if its heat of vaporization is 31.4 kJ/mol and it has a vapor pressure of mmHg at 25.0°C. 334 K

Phase Transitions Melting: change of a solid to a liquid.
H2O(s)  H2O(l) H2O(l)  H2O(s) H2O(l)  H2O(g) H2O(g)  H2O(l) H2O(s)  H2O(g) H2O(g)  H2O(s) Melting: change of a solid to a liquid. Freezing: change a liquid to a solid. Vaporization: change of a liquid to a gas. Condensation: change of a gas to a liquid. Sublimation : Change of solid to gas Deposition: Change of a gas to a solid.

Water phase changes Temperature remains constant during a phase change. Energy

Energy of Heat and Phase Change
Heat of vaporization: heat needed for the vaporization of a liquid. H2O(l) H2O(g) DH = 40.7 kJ Heat of fusion: heat needed for the melting of a solid. H2O(s) H2O(l) DH = 6.02 kJ Temperature does not change during the change from one phase to another. E.g. Start with a solution consisting of 50.0 g of H2O(s) at 0°C. Determine the heat required to heat the H2O to steam at 100.0°C.

Properties of Solutions

Classification of Matter
Solutions are homogeneous mixtures

Solute Solvent A solute is the dissolved substance in a solution.
Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Solvent A solvent is the dissolving medium in a solution. Water in salt water Water in soda

Calculations of Solution Concentration
Mass percent - the ratio of mass (in grams) of solute to mass (in grams) of solution, expressed as a percent

Molality (m) – moles of solute per kilogram of solvent

Mass/volume (m/v) % - the ratio of mass (in grams) of solute to volume of solution (in mL), expressed as a percent

Volume/volume (v/v) % - the ratio of volume (in mL) of solute to volume of solution (in mL), expressed as a percent

Mole fraction – the ratio of moles of solute to total moles of solution

Molarity (M) - the ratio of moles of solute to liters of solution

Normality (N) – moles of equilvalents/Liter of solution

Dissolving Stuff

“Like Dissolves Like” Nonpolar solutes dissolve best in nonpolar solvents Fats Benzene Steroids Hexane Waxes Toluene Polar and ionic solutes dissolve best in polar solvents Inorganic Salts Water Sugars Small alcohols Acetic acid

Heat of Solution The Heat of Solution is the amount of heat energy absorbed (endothermic) or released (exothermic) when a specific amount of solute dissolves in a solvent. Substance Heat of Solution (kJ/mol) NaOH -44.51 NH4NO3 +25.69 KNO3 +34.89 HCl -74.84

Steps in Solution Formation
H1 Expanding the solute Separating the solute into individual components H2 Expanding the solvent Overcoming intermolecular forces of the solvent molecules H3 Interaction of solute and solvent to form the solution

Enthalpy Changes in Solution
The enthalpy change of the overall process depends on H for each of these steps. Start End Start End

Why do endothermic processes sometimes occur spontaneously?
Some processes, like the dissolution of NH4NO3 in water, are spontaneous at room temperature even though heat is absorbed, not released.

Enthalpy Is Only Part of the Picture
Entropy is a measure of: Dispersal of energy in the system. Number of microstates (arrangements) in the system. b. has greater entropy,  is the favored state (more on this in chap 19)

Predicting Solution Formation
Solvent/ Solute H1 H2 H3 Hsol’n Outcome Polar/ Polar + large - large +/-small Solution forms Nonpolar + small +/- small No solution Nonpolar/ +/- small polar

Do problem 43 Do problem 45

Learning objective 2.16: The student is able to explain the properties (phase, vapor pressure, viscosity, etc.) of small and large molecular compounds.

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Learning objective 2.16: The student is able to explain the properties (phase, vapor pressure, viscosity, etc.) of small and large molecular compounds.

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