GENERAL CHEMISTRY 122 LECTURE NOTES

ORGANIC COMPOUNDS Recognition of Structures: (see handout) AlkanesKetones AlkenesCarboxylic Acids AlkynesAmines AlcoholsBenzene EthersPhenyl group Aldehydes

CHAPTER 11 LIQUIDS, SOLIDS, AND INTERMOLECULAR FORCES

PROPERTIES OF MATTER I. KINETIC MOLECULAR THEORY *All matter is composed of tiny particles (atoms or molecules) in constant motion. * Increasing temperature increases the motion of the particles. Can explain properties of matter that we observe on the macroscale on the basis of the behavior of the molecules on the nanoscale.

II. INTERMOLECULAR FORCES Forces primarily responsible for differences between (solids, liquids) and gases. What are they? 1. Attractive forces between molecules 2. They are not covalent bonds. 3. Much weaker than covalent bonds. Significant in Understanding Properties of Matter!!!

Covalent Bonds vs. Intermolecular Forces

What are they?

1.London Forces Noncovalent interaction between all molecules Due to induced dipole moments created by movement of electron clouds.

London Dispersion Forces are the only Intermolecular Force between nonpolar molecules. Factors Affecting Their Strength Greater Polarizability  Stronger London Dispersion Forces *** Larger atoms or molecules  Stronger Dispersion (London forces) (approximated to atomic or molecular mass) *** Molecular Shape  Stronger London Forces for elongated vs. compact molecules.

Which has strongest London Forces? a. Ne He Ar b. F 2 Cl 2 Br 2 I 2 Significance?

Which has the highest boiling point? (Stronger Intermolecular Forces  Higher Melting and Boiling Points) a. n-Pentane (C 5 H 12 ) vs. n-Octane (C 8 H 18 ) b. n-Octane vs. Isooctane (see next slide)

2.Dipole-Dipole Forces Noncovalent interaction between polar molecules or groups. Attractive force between a partially positive region of one molecule in close proximity to a partially negative region of another molecule The greater the polarity of a molecule, the stronger the dipole-dipole forces. Which has strongest Dipole-dipole forces? Why? H-F vs H-Br ??

Examples 1.Which types of molecules commonly have stronger intermolecular forces? Why? polar vs. nonpolar molecules 2. Which has strongest dipole-dipole forces? propane (C 3 H 8 ) or acetaldehyde (CH 3 CHO) 3. Describe / show dipole-dipole interactions for Br-Cl molecules.

3.Hydrogen Bonding Special dipole-dipole interaction. Partially positive H atom which is covalently bonded to an electronegative atom (O,N,F) has an attractive force for another electronegative atom (O,N,F). Greater # H bonds  greater intermolecular attractive forces Hydrogen Bonding is strongest of the three Intermolecular Forces !!

Examples: 1. Can they hydrogen bond? 1. Water? 2. Methanol (CH 3 OH) 3. Methane (CH 4 ) 4. Ammonia (NH 3 ) 2. Which forms greatest degree of hydrogen bonding? ethanol (CH 3 CH 2 OH) or acetic Acid (CH 3 CO 2 H)

Properties of Matter Related To Intermolecular Forces 1. Solubility Rule: “Likes Dissolve Likes” Polar and ionic compounds are soluble in polar solvents. Nonpolar compounds are soluble in nonpolar solvents. Is solid KBr more soluble in water or gasoline? Is vegetable oil more soluble in water or gasoline?

2. Greater Intermolecular Forces Related to Higher Melting and Boiling Points

3. Unusual Properties of Water **Primarily due to ability to form several “H” bonds. Water thus forms strong intermolecular forces. 1.Low mass, yet liquid at room temperature. 2.High specific heat capacity 3.High heat of vaporization, high boiling point. 4.Ice less dense than liquid water. Ice floats. Why?

PROPERTIES OF LIQUIDS RELATED TO INTERMOLECULAR FORCES I.VISCOSITY The resistance of a liquid to flow. The stronger the intermolecular forces, the greater the viscosity. Viscosity decreases as temp. increases.

II. SURFACE TENSION *Intermolecular attractions between molecules of a liquid create surface tension. (unlike a gas) * Uneveness of the forces on the surface causes the surface of the liquid to contract. * Surface tension = energy required to increase the surface area by a unit amount. * Stronger intermolecular forces  stronger or higher surface tension.

III.CAPILLARY ACTION *Molecules of liquid can interact with the molecules of the container. * Must consider attractive forces between liquid molecules compared to attractive forces between liquid and container. *What is capillary action? *What is a meniscus??

IV.VAPOR PRESSURE *Volatility – tendency of a liquid to vaporize *Vaporization- when a molecule moves from the liquid phase to the gas (vapor) phase. Why / how does this happen? (see next Figure) a. Liquid molecules have varying kinetic energy b. If kinetic energy of molecules in liquid overcomes intermolecular forces in liquid, they can escape to gas. c. Increased temp.  increased vaporization Why?

* (Equilibrium) Vapor Pressure In a closed container the liquid and gas phases of a substance come to dynamic equilibrium. What does that mean? The pressure of the gas (vapor) above a liquid at equilibium is called the equilibrium vapor pressure. What is creating the pressure? Higher volatility  more molecules in gas phase  higher vapor pressure

Example Vapor Pressure Problem Equilibrium is established between a small quantity of CCl 4 (l) and its vapor at 40 0 C in a flask having a volume of 285 mL. The total mass of vapor present is g. What is the vapor pressure of CCl 4, in mm Hg, at 40 0 C?

IV. Boiling Point of Liquid * Boiling Point- temperature at which the (equilibrium) vapor pressure equals the atmospheric pressure. *Normal BP – when atmospheric pressure = 1 atm. Higher intermolecular forces  lower vapor pressure  higher boiling point

How does size of molecules affect BP? 1. Which has highest BP? Methane(CH 4 ) Ethane(C 2 H 4 ) Propane(C 3 H 6 ) 2. Which has highest BP? Br 2 H 2 O How does change in atmospheric pressure affect BP?

PHASE CHANGES OF MATTER Phase ChangeName Solid  Liquid Liquid  Solid Melting, Fusion Freezing, Crystallization Liquid  Gas Gas  Liquid Vaporization Condensation Solid  Gas Gas  Solid Sublimation Deposition

I.PHASE TRANSITIONS a. Raindrops hit cold metal surface on car and it becomes covered with ice. b. Frozen clothes on line dry at below freezing temperatures (of H 2 O.) c. Rubbing alcohol spilled on the palm of the hand feels cool. What must be supplied or removed for phase changes?

II.Enthalpy of Phase Transitions Enthalpy (ΔH) = heat energy change under constant pressure and “temperature” conditions. If “-”  heat produced If “+”  heat required Liquid ⇌ Gas Δ H vaporization = - ΔH condensation H 2 O (l)  H 2 O (g) ΔH vap = kJ/mol H 2 O (g)  H 2 O (l) ΔH cond = kJ/mol

Solid  Liquid  Gas Endothermic processes Gas  Liquid  Solid Exothermic processes ** Consider problem solving using ΔH !!

III. Problems Using Enthalpy: Isopropyl alcohol, C 3 H 7 OH (60.0 g/mol), is used in rubbing alcohol mixtures. Alcohol on the skin cools by evaporation. How much heat is absorbed by the alcohol if 10.0 g evaporate? The enthalpy of vaporization for isopropyl alcohol is 42.1 kJ/mol.

Note: Problems with changes in temperature of a substance which also includes a phase transition. How much energy is required to heat 15g of water from –10 o C to 60 o C? Δ H fusion = kJ/mol Δ H vaporization = 40.7 kJ/mol Specific Heat Capacity solid H 2 O = 2.06 J/g o C Specific Heat Capacity liquid H 2 O = J/g o C Specific Heat Capacity gas H 2 O = 2.10 J/g o C See next Figure or Figure in text

IV.Phase Diagrams * Graphical representation to summarize conditions (pressure and temperature)under which different states of a substance are stable. See Fig *Be familiar with: a. What the lines represent b. Identify what states are present at some T and P c. Terms used

Terms Used: 1. Triple Point - P and T where all 3 phases exist in equilibrium. 2. Critical Point – endpoint of line separating liquid and gas. At this point liquid and gas are indistinguishable. Occurs at critical temperature and critical pressure. Explain 3.Supercritical Fluid - substance that exists above critical temperature and pressure. Has properties of both liquid and gas. Significance

PHASE DIAGRAM FOR WATER

PHASE DIAGRAM FOR CO 2

SOLID STATE INFORMATION I. Classified as Amorphous or Crystalline Amorphous Solids - a solid that has a disordered structure. No well defined arrangement of basic units (atoms, molecules, or ions) at nanoscale level. Examples – cement, glass, optical fibers. Crystalline Solids - a solid that has an ordered structure. Well defined symmetrical arrangement of basic units (atoms, molecules, or ions). Composed of one or more crystals with well defined 3-D structure.

II. Classification of Crystalline Solids

A. Ionic Solid – a solid that consists of positive ions (cations) and negative ions (anions) held together in a lattice by the electrical attraction of the opposite charges (ionic bonds). Very strong bonds. NaCl B. Molecular Solid – a solid that consists of atoms or molecules held together by intermolecular forces (London forces, dipole-dipole forces, and hydrogen bonds). Usually involves nonmetals. I 2, H 2 O

C. Atomic Solids 1. Metallic Solid – a solid that consists of positive cores of metal atoms held together by a surrounding sea of electrons (metallic bond). Good electrical conductors. Fe 2. Network Solid - a solid that consists of atoms held together in large networks or chains by covalent bonds. Network of nonmetal atoms. Graphite, Diamond