Chem chapter 10

Chapter 10: Intermolecular Forces: The Uniqueness of Water

  • Textbook: Chemistry (sixth edition) by Gilbert, Kirss, Bretz, Foster

  • Copyright: © 2020 W. W. Norton & Company

1. Intramolecular Forces versus Intermolecular Forces

  • Intramolecular Forces:

    • Combination of attractive and repulsive electrical forces between pairs of electrons within a molecule.

  • Intermolecular Forces:

    • Forces that act between molecules and between molecules and ions.

    • Have considerable influence on the physical properties of all substances.

2. States of Matter

  • 3 States:

    • (a) Solid

    • (b) Liquid

    • (c) Gas

3. Boiling Points and Atomic Number

  • The boiling point increases as the atomic number increases in pure compounds.

  • Separating liquid-phase particles that are attracted to each other requires energy.

    • The stronger the particles’ attractions for each other, the greater the energy needed to separate them.

    • As a result, higher boiling points are observed with increased energy requirements.

4. Interactions Involving Polar Molecules: Ion-Dipole Interaction

  • Ion–Dipole Interaction:

    • Attractive force between an ion and a molecule with a permanent dipole.

    • Example:

    • Na+ attracts O in H2O

    • Cl- attracts H in H2O

5. Solvation of Ions

  • When a cation exists in solution, it is surrounded by the negative dipole ends of water molecules.

  • When an anion exists in solution, it is surrounded by the positive dipole ends of water molecules.

6. Sphere of Hydration (Solvation)

  • Defined as the cluster of water molecules surrounding an ion in aqueous solution.

  • Inner Sphere of Hydration

  • Outer Sphere of Hydration

  • Related Interactions:

    • Ion-dipole interaction

    • Dipole-dipole interaction

7. Interactions of Nonpolar Molecules

  • Dispersion (London) Forces:

    • Intermolecular force found between nonpolar molecules caused by the presence of temporary dipoles.

  • Temporary (Induced) Dipole:

    • Separation of charge produced in an atom or molecule by an uneven distribution of electrons.

  • Polarizability:

    • The relative ease with which the electron cloud in a molecule, ion, or atom can be distorted, inducing a temporary dipole.

7a. Induced Dipole

  • Described as a momentary shift in electron density, creating dispersion forces.

7b. London Dispersion Forces

  • Exist between all molecules and are a function of molecular polarizability.

    • Polarizability Measurement:

    • Determined by the ease of distortion of the electron cloud.

    • Larger molecules with more electrons show greater polarizability.

    • Increased surface area also contributes to higher dispersion forces.

    • London dispersion forces increase with molecular weight.

7c. Strength of London Dispersion Forces

  • Factors Affecting Strength:

    1. Size of Atoms/Molecules:

    • Larger atoms/molecules are more polarizable.

    1. Shape of Molecules:

    • Increased surface area = increased interactions.

    • Linear molecules have higher dispersion than branched molecules of similar molecular weight.

8. Boiling Point and Intermolecular Forces

  • Boiling point reflects the strength of intermolecular forces, which increase with polarization of bonds.

  • Boiling point also increases with molecular weight and surface area.

8a. Effect of Size on Dispersion

  • Table 10.2: Boiling Points of the Halogens

    • F₂: Molar Mass = 38 g/mol, Boiling Point = 85 K

    • Cl₂: Molar Mass = 71 g/mol, Boiling Point = 239 K

    • Br₂: Molar Mass = 160 g/mol, Boiling Point = 332 K

    • I₂: Molar Mass = 254 g/mol, Boiling Point = 457 K

8b. Effect of Shape on Dispersion

  • Zigzag (straight-chain) hydrocarbons have more surface area, resulting in stronger London forces and higher boiling points.

  • Branched Hydrocarbons:

    • Cannot stack tightly together, leading to lower boiling points.

9. Dipole/Induced Dipole Interactions

  • Polar molecules (e.g., water) can create or induce dipoles in nonpolar molecules.

  • Nonpolar solids like I₂ are insoluble in water.

  • Nonpolar gases (O2, N2, H2, CO2) show limited solubility in water, despite the necessity of oxygen in aquatic environments.

10. Dipole-Induced Dipole Interaction

  • Proximity of polar molecules can induce a dipole in nonpolar molecules, resulting in temporary bonded interactions.

11. Induced Dipole Forces

  • Measured by polarizability; larger molecules are more easily polarized.

  • Looser-held electrons relate to higher polarizability, aligning with trends in solubility.

12. Dipole-Dipole Interactions: Hydrogen Bonding

  • Dipole-Dipole Interaction:

    • Attractive force between polar molecules.

  • Hydrogen Bond:

    • Strongest dipole-dipole interaction.

    • Conditions for existence:

    • High electronegativity

    • Presence of lone pairs

    • H interacting with N, O, or F

    • Energy range: 4extto25extkJ/mol4 ext{ to } 25 ext{ kJ/mol}.

    • Crucial in biological systems (e.g., DNA, RNA, protein structures).

13. Hydrogen Bonding Phenomena

  • Water Molecule Network:

    • Hydrogen bonding results in unique properties like capillary action, surface tension, and the flotation of ice.

14. Boiling Points of Simple Hydrogen Compounds

  • Notably high boiling points observed for water, ammonia, and HF relative to their molecular weights:

    • Water: 18 g/mol

    • Ammonia: 20 g/mol

    • HF: 17 g/mol

15. Ice Structure

  • Ice exhibits an open lattice structure, making it less dense than liquid water, allowing it to float on water's surface.

16. Hydrogen Bonding in Biological Structures

  • H-bonding between complementary strands of DNA crucial for structure:

    • Adenine (A) and Thymine (T) form two hydrogen bonds.

    • Guanine (G) and Cytosine (C) form three hydrogen bonds.

17. Relative Strengths of Intermolecular Forces

  • Table 10.3: Type of Force and Phenomena Explained

    • Ion-dipole: Moderate strength (10-50 kJ/mol) - Example: NaCl in water

    • Hydrogen bonding: Moderate (10-40 kJ/mol) - Example: Water expands when it freezes

    • Dipole-dipole: Weak (3-4 kJ/mol) - Example: Formaldehyde boiling point.

    • Dispersion (London Forces): Weak (1-10 kJ/mol) - Included in all molecular interactions.

18. Summary of Intermolecular Forces

Force Type

Strength

Description

Ion-Dipole

Moderate (10-50 kJ/mol)

Occurs between ions and polar solvents

Hydrogen Bond

Moderate (10-40 kJ/mol)

Occurs between molecules with O-H, N-H, F-H bonds

Dipole-Dipole

Weak (3-4 kJ/mol)

Occurs between polar molecules

London Dispersion

Weak (1-10 kJ/mol)

Occurs between ALL molecules

19. Practice Questions

  • What is the strongest intermolecular force between ethanol molecules?

    • A. Dispersion forces

    • B. Dipole-dipole forces

    • C. Hydrogen bonds

    • D. Ion-dipole forces

  • What intermolecular forces exist between dimethyl ether molecules?

    • A. Dispersion forces only

    • B. Dipole-dipole forces only

    • C. Dispersion forces and dipole-dipole forces

    • D. Dispersion forces, dipole-dipole forces, and hydrogen bonds

20. Practice: Predicting Boiling Points

  • Rank the following compounds in order of increasing boiling point: CH3OH, CH3CH2CH2CH3, and CH3CH2OCH3

20a. Collect and Organize

  • Given molecular formulas to rank based on boiling points.

20b. Analyze

  • Identify main intermolecular forces for each compound and rank based on their strengths.

20c. Solve

  • Lowest: CH3CH2CH2CH3 - Only dispersion forces, weakest.

  • Middle: CH3CH2OCH3 - Contains O (dipole) alongside dispersion forces.

  • Highest: CH3OH - Capable of hydrogen bonding, strongest of the intermolecular forces.

21. Summary of Predictions on Boiling Points

  • The boiling points of substances increase with the strength of interactive forces.

22. Summary of Intermolecular Forces Interaction

  • Are polar molecules involved?

  • Are ions involved?

  • Are polar molecules involved?

  • Are Hydrogen atoms bonded to O, N or F involved?

    • London Forces (induced dipoles): No

    • Dipole-Dipole forces: Yes

    • Hydrogen Bonding: Yes

    • Ion-Dipole forces: Yes

23. Relating Properties of Solvents and Solubility

  • Stronger intermolecular forces lead to lower vapor pressure and higher boiling points.

24. Vapor Pressure

  • Definition: Pressure exerted by molecules of a substance in the gas phase in equilibrium with its liquid phase.

  • Factors Affecting Vapor Pressure:

    1. Temperature: Higher temps increase kinetic energy, enhancing vapor pressure.

    2. Surface Area: More molecules at the surface increase vapor pressure.

    3. Intermolecular Forces: Stronger forces require more energy to escape, thus reducing vapor pressure.

25. Enthalpy of Vaporization

  • The pressure exerted by vapor molecules above a liquid in a closed container equates to equilibrium vapor pressure, where the rate of evaporation equals the rate of condensation.

26. Relative Vapor Pressures

  • Characteristics:

    • Strong intermolecular forces lead to: low vapor pressure, high boiling point, low evaporation rates.

    • Weak intermolecular forces yield: high vapor pressure, low boiling point, high evaporation rates.

27. Vapor Pressure vs. Temperature

  • Graph: Vapor pressure plotted against temperature reveals varying boiling points for substances (including water, ethanol, and diethyl ether).

28. Clausius–Clapeyron Equation

  • The fundamental equation relating vapor pressure to temperature and enthalpy changes.

  • Equation: ln(P<em>vap,T2)−ln(P</em>vap,T1)=−ΔH<em>vapR(1T</em>2−1T1)ln(P<em>{vap, T2}) - ln(P</em>{vap, T1}) = -\frac{\Delta H<em>{vap}}{R} (\frac{1}{T</em>2} - \frac{1}{T_1})

29. Practice: Calculating Vapor Pressure

  • Scenario: Calculate vapor pressure for water in a panting dog's breath at 35 °C

  • Given: Enthalpy of Vaporization = 40.7 kJ/mol

29a. Collect and Organize

  • Required calculations based on Henry's Law and given conditions.

29b. Analyze

  • Necessary conversions for calculation purposes.

  • Water vapor pressure at 20 °C = 24 torr; expect this to rise around 50 torr at 35 °C.

29c. Solve

  • Calculation steps and outcomes.

30. Phase Diagrams

  • Definition: Graphical representation of the physical states of a substance concerning temperature and pressure.

  • Equilibrium Lines: Represent phase changes (solid/liquid, liquid/gas, solid/gas).

  • Supercritical Region: Unique phase of matter with properties intermediate to gas and liquid.

31. Key Features of Phase Diagrams

  • Triple Point: Conditions for all three phases coexisting.

  • Critical Point: Specific temperature and pressure where liquid and gas phases have the same density.

31a. Critical Temperature and Pressure

  • Critical Temperature: Highest temperature allowing liquid existence.

  • Critical Pressure: The pressure needed to liquefy a substance at its critical temperature.

31b. Supercritical Fluids

  • Unique properties, utilized in caffeine extraction without losing flavor.

32. Properties of Water

  • Due to intermolecular forces, water has diverse properties such as high surface tension and unique density.

  • Density of Ice: Lattice structure of ice enables it to be less dense than liquid water, crucial for aquatic life.

33. Solubility: Like Dissolves Like

  • Solubility is determined by relative strengths of solute-solvent attractions; polar solutes dissolve in polar solvents.

33a. Solubility Rules

  • Soluble Compounds: Group I ions, certain anions like NO3– and acetates.

  • Insoluble Compounds: Most hydroxides, sulfides, carbonates, and phosphates except with Group I/NH4+ compounds.

34. Predicting Solubility in Water

  • Evaluate compounds to predict solubility based on their molecular structure and polarity.

34a. Practice: Analyze Four Compounds

  • Determining solubility concludes that polar molecules (NH3 and HF) are soluble while nonpolar ones (CCl4 and N2) exhibit limited solubility.

35. Combination of Intermolecular Forces

  • Hydrophobic vs Hydrophilic:

    • Hydrophobic: diminishes water solubility

    • Hydrophilic: promotes water solubility

35a. Surfactants

  • Molecules exhibiting both hydrophobic and hydrophilic nature, allowing them to interact with both polar and nonpolar substances.

36. Insights on Gas Solubility

  • Solubility decreases with an increase in temperature due to intensified kinetic energy overpowering solute-solvent interactions.

37. Henry’s Law

  • Formula that quantifies gas solubility in liquids in relation to gas pressure: C<em>gas=k</em>HP<em>gasC<em>{gas} = k</em>H P<em>{gas} where $C{gas}$ denotes concentration, $kH$ is the Henry's Law constant, and $P{gas}$ is the partial pressure surrounding the solvent.

37a. Practice: Solve Gas Concentration in Beverages

  • Calculation steps to find concentration of carbon dioxide in soft drinks utilizing Henry’s Law constant.

38. Summary of Henry's Law Calculations

  • Norm for determining gas concentration in solutions relying on provided constants and pressures.

  • Emphasizes the impact of molecular structure on solubility and interactions.