Unit 6: States of Matter, Intermolecular Forces, Solutions

Intramolecular Forces vs. Intermolecular Forces Intramolecular Forces

o Within the moleculeo Molecules are formed by sharing electrons between the atoms

Intermolecular Forceso Between moleculeso Dipole-Dipole

Hydrogen Bondingo London Dispersion

Intramolecular forces are stronger than intermolecular forces! Dotted lines are the IMFs between the water molecules

Phase Changeso When a substance changes from solid to liquid to gas, the molecules must remain intact.o The changes in state are due to changes in the forces among molecules rather than in those within

the molecules.o When phases change, IMFs are overcome.

Intermolecular Forces aka Van der Waal Forces London Dispersion Forces

o Instantaneous dipole that occurs accidentally in a given atom induces a similar dipole in a neighboring atoms.

o Significant in large atoms/moleculeso Occurs in all molecules, including nonpolar

Primary force in nonpolar molecules

Dipole-Dipole Forceso Dipole moment

molecules with polar bonds often behave in an electric field as if they had a center of positive charge and a center of negative charge.

Molecules can attract each other electrostatically, lining up so that the positive and negative ends are close to each other

Only ~1% as strong as covalent or ionic bondso Hydrogen Bonding

Strong dipole-dipole forces Hydrogen is bound to a highly electronegative atom (nitrogen, oxygen, fluorine)

Hydrogen is then electrostatically attracted to a lone pair on the electronegative atom or adjacent molecules

Melting and Boiling Points In general, the stronger the intermolecular forces, the higher the melting and boiling points The Boiling Points of the Covalent Hydrides of the Elements in Groups 4A, 5A, 6A, and 7A

1. Which molecule is capable of forming stronger intermolecular forces? Explain.

N2 H2O

2. Draw two Lewis structures for the formula C2H6O and compare the boiling points of the two molecules.

3. Which gas would behave more ideally at the same conditions of P and T? Why?CO or N2

4. Indicate all the various types of intermolecular attractive forces that may operate in each of the following:a. CH3OH (l)b. Xe (l)c. H2S (l)d. ClF (l)e. Ca(NO3)2 (s)

5. For each of the following substances, predict which will have the higher melting point and indicate why:a. CuBr2, Br2

b. CO2, SiO2

c. S, Crd. CsBr, CaF2

Liquids Low compressibility, lack of rigidity, and high density compared with gases. Surface tension – resistance of a liquid to an increase in its surface area:

o Liquids with large intermolecular forces tend to have high surface tensions. Capillary action – spontaneous rising of a liquid in a narrow tube:

o Cohesive forces – intermolecular forces among the molecules of the liquid.o Adhesive forces – forces between the liquid molecules and their container.

Viscosity – measure of a liquid’s resistance to flow:o Liquids with large intermolecular forces or molecular complexity tend to be highly viscous.

6. Which force dominates alongside the glass tube – cohesive or adhesive forces?

Convex Meniscus Formed Convex Meniscus Formed by Nonpolar Liquid Mercury by Nonpolar Liquid Mercury

HW 6A

Atomic Molecular IonicCovalent Network Metallic Lone Atoms Nonmetals Salts

Examples

Type of Bonding or IMF

Typical Properties

Solids Amorphous Solids:

o Disorder in the structureso Glass

Crystalline Solids:o Ordered Structureso Unit Cellso Examples of Three Types of Crystalline Solids

7. Indicate the type of crystal (lone atom, molecular, metallic, colvalent or ionic)

a. O2

b. H2Sc. Agd. KCle. Sif. Al2(SO4)3

g. Neh. SiO2

i. NH3

j. MgOk. NaOHl. CH4

Bonding Models for Metals Electron Sea Model - A regular array of cations in a “sea” of mobile valence electrons.

Band Model (MO Model) - Electrons are assumed to travel around the metal crystal in molecular orbitals formed from the valence atomic orbitals of the metal atoms.

Two Types of Alloys Substitutional alloy - some of the host metal atoms are replaced by other

metal atoms of similar size.o Brass

Interstitial alloy - some of the holes in the closest packed metal structure are occupied by small atoms.

o Steel

Network Solids

Diamond Graphite Ceramics

o Typically made from clays (which contain silicates) and hardened by firing at high temperatures.o Nonmetallic materials that are strong, brittle, and resistant to heat and attack by chemicals.

Semiconductorso n-type semiconductor – substance whose conductivity is increased by doping it with atoms having

more valence electrons than the atoms in the host crystal.o p-type semiconductor – substance whose conductivity is increased by doping it with atoms having

fewer valence electrons than the atoms of the host crystal.

Ionic Solids stable, high melting substances held together by the strong electrostatic forces that exist between oppositely

charged ions.Molecular Solids

less stable, lower melting substances held together by either polar or nonpolar covalent bonds

HW 6B

Vapor Pressure Pressure due to particles of a substance in the vapor phase above its liquid in a closed container at a given

temperature. The weaker the forces holding the liquid together, the higher the vapor pressure of the liquid will be.

Behavior of a Liquid in a Closed Container

Initially Equilibrium

Pressure of the vapor present at equilibrium The system is at equilibrium when no net change occurs in the amount of liquid or vapor because the two

opposite processes exactly balance each other.

8. What is the vapor pressure of water at 100°C? How do you know?

Liquids in which the intermolecular forces are large have relatively low vapor pressures. Vapor pressure increases significantly with temperature.

Clausius–Clapeyron EquationPvap = vapor pressureΔHvap = enthalpy of vaporizationR = 8.3145 J/K·molT = temperature (in kelvin)

1

2

vap, vap

vap, 2 1

1 1ln =

T

T

P HP R T T

9. The vapor pressure of water at 25°C is 23.8 torr, and the heat of vaporization of water at 25°C is 43.9 kJ/mol. Calculate the vapor pressure of water at 65°C.

Heating Curve for Water

10. Which would you predict should be larger for a given substance: ΔHvap or ΔHfus? Explain

HW 6C

Phase Diagrams A convenient way of representing the phases of a substance as a function of temperature and pressure:

o Triple pointo Critical pointo Phase equilibrium lines

11. As intermolecular forces increase, what happens to each of the following? Why?a. Boiling pointb. Viscosityc. Surface tensiond. Enthalpy of fusione. Freezing pointf. Vapor pressureg. Heat of vaporization

HW 6C

Solutions

Solution Composition

12. You have 1.00 mol of sugar in 125.0 mL of solution. Calculate the concentration.

13. You have a 10.0 M sugar solution. What volume of this solution do you need to have 2.00 mol of sugar?

14. Consider the separate solutions of NaOH and KCl made by dissolving 100.0 g of each solute in 250.0 mL of solution. Calculate the concentration of each solution.

15. What is the percent by mass concentration of glucose in a solution made by dissolving 5.5 g of glucose in 78.2 g of water?

16. A solution of phosphoric acid was made by dissolving 8.00 g of H3PO4 in 100.0 mL of water. Calculate the mole fraction of H3PO4. Assume water has a density of 1.00 g/mL.

Formation of a Liquid Solution: THE DISSOLVING PROCESS Separating the solute into its individual components (expanding the solute). Overcoming IMFs in the solvent to make room for the solute (expanding the solvent). Allowing the solute and solvent to interact to form the solution. Steps in the dissolving process

o Steps 1 and 2 require energy since forces must be overcome to expand the solute and solvent.o Step 3 usually releases energy.o Steps 1 and 2 are endothermic, and step 3 is often exothermic.o Enthalpy change associated with the formation of the solution is the sum of the ΔH values for the

steps:ΔHsoln = ΔH1 + ΔH2 + ΔHs3

o ΔHsoln may have a positive sign (energy absorbed) or a negative sign (energy released)

Enthalpy (Heat) of Solution

17. Explain why water and oil do not mix. In your explanation, be sure to address how ΔH plays a role.

In General: One factor that favors a process is an increase in probability of the state when the solute and solvent are

mixed.

Processes that require large amounts of energy tend not to occur.

Overall, just remember that “like dissolves like”…but don’t say that as an explanation!!

Factors Effecting Solubility Structure effects: Polarity

o Hydrophobic – nonpolar substanceso Hydrophilic – polar substances

Pressure effects: Henry’s Lawo Little effect on solubility if solids and liquidso Henry’s Law

C=kPC = concentration of dissolved gask = constantP = partial pressure of gas solute above the solution

o Amount of gas dissolved in a solution is directly proportional to the pressure of the gas above the solution.

Temperature effects: affecting aqueous solutions Solubilities of Solids Solubilities of Gases

o Although the solubility of most solids in water increases with temperature, the solubilities of some substances decrease with increasing temperature.

o Predicting temperature dependence of solubility is very difficult.o Solubility of a gas in solvent typically decreases with increasing temperature.

HW 6D

Colligative Properties Properties that depend on the number of molecules or ions of solute present and not on what the particles

are. Vapor Pressure Lowering Boiling Point Elevation Freezing Point Depression

18. If you have 1 molar glucose in 1 L of water and 1 molar NaCl in 1 liter of water, which solute has a greater effect on any colligative property?

19. If you have 1 molar Na3PO4 in 1 L of water and 1 molar NaCl in 1 liter of water, which solute has a greater effect on any colligative property?

Real Life Applicationso Salt on icy roads

o Antifreeze in radiator

o Salt in water when cooking

o Higher altitude when cooking

HW 6E