Intermolecular Forces

Chapter 6 – Intermolecular Forces

Definitions

  • Intramolecular forces: Forces within a molecule, comprising electrostatic/ionic or covalent forces (bonds).

  • Intermolecular forces (IMFs): Forces that act between one molecule and another, also referred to as "Van der Waals forces." These forces hold molecules together in a clump.

  • Cohesion: The forces that bind molecules together to form a bulk material.

Phases of Matter

  • Liquid:

    • Strong interactions between neighboring molecules.

    • High degree of order within the structure.

  • Solid:

    • Moderately strong interactions between neighboring molecules.

    • Intermediate degree of order.

  • Gas:

    • Weak interactions between neighboring molecules.

    • Low degree of order.

Energy Forms in Matter

  • All matter, whether solid, liquid, or gas, has two counteracting forms of energy:

    • Potential Energy: Energy that tends to give order to matter, holding matter together.

    • Kinetic Energy: Energy of molecular motion, tending to disorganize matter.

  • Measure of IMFs: The boiling point of a substance is the best measure of the strength of intermolecular forces. A higher boiling point indicates stronger IMFs.

Types of Intermolecular Forces

  • Weak Electrostatic Forces: Intermolecular forces generally described as weak and of varying strength between molecules.

Types of Intermolecular Forces Explained

  • Dipole-Dipole Forces:

    • Occur when two polar molecules attract each other.

  • Ion-Dipole Forces:

    • A polar molecule attracts an ion.

  • Hydrogen Bonding:

    • This occurs when nitrogen (N), oxygen (O), or fluorine (F) with a negative charge attracts hydrogen in the configuration of H-N, H-O, or H-F which has a positive charge.

  • London Dispersion Forces (LDFs):

    • Present in any molecule with electrons. More electrons or greater molecular weight correlates to greater LDFs.

Cumulative Nature of IMFs

  • Intermolecular forces are cumulative between two molecules:

    • Energy Values for Various Bonds (typical NRG in kJ/mol):

    • Covalent Bond: 250 kJ/mol

    • Ionic Bond: 250 kJ/mol

    • Dipole-Dipole: 5 - 25 kJ/mol

    • London Dispersion: 0.5 - 40 kJ/mol

    • Hydrogen Bond: 10 - 40 kJ/mol

London Dispersion Forces Diagram

  • Electrostatic Attraction Example:

    • (a) Helium atom 1

    • (b) Helium atom 2

Practice Problems

  • (Hydrogen Bonding) Identify which can form hydrogen bonds with itself:

    • H3C-OH

    • H3C-F

    • H3C-Cl

    • H3C-NH2

  • Identify hydrogen bond formation with a molecule of N(CH3)3:

  • Determine hydrogen bond formation with a molecule of H3C-CH3:

    • No hydrogen bond formation occurs due to the lack of electronegative atoms capable of forming such bonds with hydrogen.

    • Where X and Y are either a nitrogen (N), oxygen (O), or fluorine (F) donor and acceptor.

  • Compare boiling points:

    • propane (C3H8, M = 44 g/mol) vs. butane (C4H10, M = 58 g/mol)

    • methane (CH4, M = 16 g/mol) vs. chloromethane (CH3Cl, M = 50 g/mol)

    • chloromethane (CH3Cl, M = 50 g/mol) vs. methanol (CH3OH, M = 32 g/mol)

    • viscosity comparison at the same temperature of chloromethane and methanol.

Solutions and Terms Related to Solvation

  • Solvent: The component present in greater quantity.

  • Solute: Other substances that are dissolved in the solvent.

  • Solvation: An ion or a molecule is surrounded by solvent molecules.

  • Many solutes can dissolve in a solvent, exemplified by beer.

Principles of Solubility

  • "Like dissolves like": The more chemically similar two compounds are, the more likely it is that one is soluble in the other.

    • Polar compounds dissolve polar compounds.

    • Non-polar compounds dissolve non-polar compounds.

  • Hydrophilic: "water-attracting" (polar) substances.

  • Hydrophobic: "water-repelling" (non-polar) substances.

Comparisons of Compound Solubility

Hydrophilic

Hydrophobic

Water

-

-

Non-polar Compounds

-

Oil

Polar Compounds

-

Salt/Ions, Grease/Fats

Small Alcohols (ROH)

Sugars (sucrose)

Benzene, C2Cl4, CCl4

Hydrocarbons (hexane)

-

-

Surfactants (SOAP)

-

-

Additional Practice Questions on Solubility

  • Determination in water:

    • Which dissolves better, H3C-OH or H3C-I? Explain.

  • Determination in benzene:

    • Which dissolves better, H3C-F or H3C-CH3? Explain.

Concepts of Gas Pressure

  • Pressure: Defined as Force/Area (similar to atmospheric pressure).

  • Conversions:

    • 760 mm Hg = 760 torr = 1 atm

    • SI unit for pressure = Pascal (Pa), where 1 Pa = 1 kg/m·s².

    • Exact conversion: 1 atm = 101.325 kPa.

  • STP (Standard Temperature and Pressure): Defined as 1 atm and 273 K.

  • States of matter can be influenced by altering temperature and/or pressure.

Phase Changes and Properties

  • Normal Boiling Point (Tb): Temperature at which vapor pressure equals 1 atm.

  • Normal Melting Point (Tm): The temperature at which a liquid freezes at 1 atm.

  • Phase Boundaries: Lines that separate different states of matter on a phase diagram; points on these boundaries indicate conditions where two phases coexist in dynamic equilibrium.

  • Triple Point: The specific conditions where three phase boundaries meet, allowing all three phases to coexist in dynamic equilibrium.

Critical Properties of Substances

  • Critical Temperature (Tc): The temperature above which a substance cannot be condensed to a liquid.

  • Critical Pressure (Pc): The minimum pressure at which a substance can exist in its gaseous state.

  • Supercritical Fluid: A state of matter that exists above its critical temperature and pressure.

Problems on Phases and Critical Points

  • Determine the phase at specific conditions:

    • At 0.5 atm and 100 K.

    • Find normal boiling and melting points.

    • Find the temperature and pressure at the triple point.

    • Determine the temperature and pressure at the critical point.

  • Determine the phase at 88 atm and 300 °C.

  • Find the temperature and pressure at the critical point and triple point.

Chapter 6 – Intermolecular Forces
Definitions

  • Intramolecular forces: Forces within a molecule, comprising electrostatic/ionic or covalent forces (bonds). These are strong forces that determine the identity and chemical properties of a compound.

  • Intermolecular forces (IMFs): Forces that act between one molecule and another, also referred to as "Van der Waals forces." These forces are much weaker than intramolecular forces but are responsible for the physical properties of substances, such as boiling point, melting point, and viscosity. They hold molecules together in a clump in condensed phases (liquids and solids).

  • Cohesion: The forces that bind molecules of the same substance together to form a bulk material. Adhesion refers to forces between different substances.

Phases of Matter

The physical state of a substance (solid, liquid, or gas) is determined by the balance between the kinetic energy of its molecules and the strength of the intermolecular forces holding them together.

  • Solid:

    • Strongest interactions between neighboring molecules due to tightly packed arrangement.

    • High degree of order; molecules are held in fixed positions and can only vibrate around these positions.

    • Possesses a definite volume and definite shape.

  • Liquid:

    • Moderately strong interactions between neighboring molecules, allowing them to remain close but also to move past one another.

    • Intermediate degree of order; molecules are randomly arranged but still in contact.

    • Possesses a definite volume but an indefinite shape, taking the shape of its container.

  • Gas:

    • Weakest interactions between neighboring molecules; molecules are far apart and move rapidly and randomly.

    • Low degree of order; molecules are essentially independent of each other.

    • Possesses an indefinite volume and indefinite shape, expanding to fill its container.

Energy Forms in Matter

  • All matter, whether solid, liquid, or gas, has two counteracting forms of energy:

    • Potential Energy: Energy associated with the attractive forces (IMFs) between molecules, which tends to pull molecules together and give order to matter.

    • Kinetic Energy: Energy of molecular motion (vibration, rotation, translation), which tends to spread molecules apart and disorganize matter. As temperature increases, kinetic energy increases.

  • Measure of IMFs: The boiling point of a substance is the best macroscopic measure of the strength of intermolecular forces. A higher boiling point indicates stronger IMFs because more kinetic energy (heat) is required to overcome these attractive forces and allow molecules to escape into the gas phase.

Types of Intermolecular Forces

  • Weak Electrostatic Forces: Intermolecular forces are generally described as weak because they involve partial charges or temporary dipoles, unlike the full charges or shared electron pairs in intramolecular bonds. Their strength varies considerably between different types.

Types of Intermolecular Forces Explained

  • Dipole-Dipole Forces:

    • Occur between two polar molecules. Polar molecules possess a permanent dipole moment due to unequal sharing of electrons and an asymmetrical molecular geometry.

    • The positive end of one polar molecule is electrostatically attracted to the negative end of an adjacent polar molecule.

    • The strength of dipole-dipole forces depends on the magnitude of the dipole moment of the molecules (i.e., how polar they are).

  • Ion-Dipole Forces:

    • Occur when a polar molecule attracts an ion (either cation or anion).

    • These are typically the strongest type of IMF, crucial for the dissolution of ionic compounds in polar solvents (e.g., NaCl dissolving in water). The charged ion is attracted to the opposite partial charge of the polar molecule.

  • Hydrogen Bonding:

    • A particularly strong type of dipole-dipole interaction. It occurs when a hydrogen atom (H) is covalently bonded to a highly electronegative atom — nitrogen (N), oxygen (O), or fluorine (F) — and this H atom is attracted to a lone pair of electrons on another N, O, or F atom in an adjacent molecule.

    • The high electronegativity of N, O, or F pulls electron density away from H, giving H a significant partial positive charge (\delta+). The small size of H allows for a close approach to the lone pair on the acceptor atom.

    • Hydrogen bonding is vital for the properties of water, DNA structure, and protein folding.

  • London Dispersion Forces (LDFs):

    • Also known as induced dipole-induced dipole forces, or Van der Waals forces in a broad sense.

    • Present in all molecules, polar and nonpolar, and are the only IMFs present in nonpolar molecules and noble gas atoms.

    • Arise from temporary, instantaneous dipoles created by the random movement of electrons around the nucleus. At any given moment, electron distribution can be asymmetrical, creating a transient dipole that can induce a dipole in a neighboring molecule.

    • The strength of LDFs increases with:

      • Number of electrons (or molar mass): Larger atoms/molecules have more electrons, leading to larger and more easily distorted electron clouds (higher polarizability), resulting in stronger instantaneous dipoles.

      • Molecular shape: Molecules with larger surface areas or elongated shapes allow for more points of contact between molecules, leading to stronger LDFs (e.g., n-pentane has stronger LDFs than neopentane).

Cumulative Nature of IMFs

  • Intermolecular forces are cumulative; while individual forces are weak compared to covalent or ionic bonds, their combined effect significantly influences a substance's macroscopic properties.

  • Energy Values for Various Bonds (typical NRG in kJ/mol):

    • Covalent Bond: 250 - 1000 kJ/mol (strongest intramolecular forces)

    • Ionic Bond: 250 - 4000 kJ/mol (strongest intramolecular forces)

    • Dipole-Dipole: 5 - 25 kJ/mol (weakest type of IMF, but stronger than LDFs for similar size)

    • London Dispersion: 0.5 - 40 kJ/mol (present in all molecules, can be significant for large molecules)

    • Hydrogen Bond: 10 - 40 kJ/mol (strongest type of IMF)

London Dispersion Forces Diagram

  • Electrostatic Attraction Example: This diagram typically illustrates how instantaneous dipoles lead to attraction:

    • (a) Helium atom 1: At a given instant, electron density may be uneven, creating a temporary dipole (one side \delta- , other side \delta+ ).

    • (b) Helium atom 2: The temporary dipole in atom 1 induces a complementary temporary dipole in atom 2, leading to a weak electrostatic attraction between them.

Practice Problems

  • (Hydrogen Bonding) Identify which can form hydrogen bonds with itself: (For H-bonding with itself, the molecule must contain H bonded to N, O, or F, and have a lone pair on N, O, or F to act as an acceptor). These molecules can act as both a donor and acceptor.

    • H3C-OH (Methanol): Yes, O is bonded to H, and O has lone pairs.

    • H3C-F (Fluoromethane): No, H is not bonded to F. F is electronegative but doesn't have an H directly attached to it that can hydrogen bond from the molecule.

    • H3C-Cl (Chloromethane): No, H is not bonded to N, O, or F. Cl is not electronegative enough to participate in hydrogen bonding.

    • H3C-NH2 (Methylamine): Yes, N is bonded to H, and N has a lone pair.

  • Identify hydrogen bond formation with a molecule of N(CH3)3 (trimethylamine): N(CH3)3 can accept hydrogen bonds because the N atom has a lone pair. However, it cannot donate a hydrogen bond because it lacks H atoms directly bonded to N, O, or F.

  • Determine hydrogen bond formation with a molecule of H3C-CH3 (ethane):

    • No hydrogen bond formation occurs because there are no electronegative atoms (N, O, F) capable of forming such bonds with hydrogen, nor are there any H atoms bonded to N, O, or F which can act as a donor.

    • Where X and Y are either a nitrogen (N), oxygen (O), or fluorine (F) donor and acceptor and form the H-bond (X-H \cdots Y).

  • Compare boiling points (higher boiling point indicates stronger IMFs):

    • Propane (C3H8, M = 44 g/mol) vs. butane (C4H10, M = 58 g/mol): Butane has a higher boiling point. Both are nonpolar and only have LDFs. Butane has more electrons and a larger molar mass, leading to stronger LDFs.

    • Methane (CH4, M = 16 g/mol) vs. chloromethane (CH3Cl, M = 50 g/mol): Chloromethane has a higher boiling point. Methane is nonpolar (LDFs only). Chloromethane is polar (LDFs + dipole-dipole forces). The dipole-dipole forces in CH3Cl are stronger than the LDF increase due to higher mass in this case.

    • Chloromethane (CH3Cl, M = 50 g/mol) vs. methanol (CH3OH, M = 32 g/mol): Methanol has a significantly higher boiling point. Chloromethane has LDFs and dipole-dipole forces. Methanol has LDFs, dipole-dipole forces, and can form hydrogen bonds (due to O-H bond), which are the strongest IMFs present here.

    • Viscosity comparison at the same temperature of chloromethane and methanol: Methanol would have higher viscosity. Viscosity is a measure of a fluid's resistance to flow, which increases with stronger IMFs because molecules are more resistant to moving past each other. Methanol's hydrogen bonding contributes to much stronger IMFs than chloromethane's dipole-dipole forces.

Solutions and Terms Related to Solvation

  • Solvent: The component present in greater quantity in a solution, which dissolves the solute.

  • Solute: Other substances that are dissolved in the solvent, typically present in lesser quantity.

  • Solvation: The process by which solvent molecules surround and interact with an ion or a molecule of the solute. When the solvent is water, this process is called hydration.

  • Many solutes can dissolve in a solvent, exemplified by beer (a solution of various solutes in water).

Principles of Solubility

  • "Like dissolves like": This fundamental principle states that the more chemically similar two compounds are, particularly in terms of their polarity and ability to form specific types of IMFs, the more likely it is that one is soluble in the other. For a solute to dissolve in a solvent, the attractive forces between solute and solvent molecules must be comparable to or stronger than the forces holding the solute molecules together and the solvent molecules together.

    • Polar compounds dissolve polar compounds. This is because polar-polar interactions (e.g., dipole-dipole, hydrogen bonding) can overcome the existing polar-polar interactions and form new ones.

    • Non-polar compounds dissolve non-polar compounds. This is because nonpolar compounds primarily interact via LDFs, and mixing them allows for similar LDFs to form between the molecules.

  • Hydrophilic: "water-attracting" substances. These are typically polar molecules or ions that can form strong IMFs (especially hydrogen bonds or ion-dipole forces) with water.

  • Hydrophobic: "water-repelling" substances. These are typically non-polar molecules that cannot form significant attractive forces with water molecules, leading to their exclusion from the water.

Comparisons of Compound Solubility

  • Determination in Water:

    • Which dissolves better, H3C-OH (Methanol) or H3C-I (Iodomethane)? Explain.

      • H3C-OH will dissolve better in water. Methanol is polar and can form hydrogen bonds with water due to its -OH group. Iodomethane is also polar but cannot form hydrogen bonds with water, and its primary IMFs with water would be weaker dipole-dipole and LDFs. The strong H-bonding drives methanol's solubility.

  • Determination in Benzene (C6H6, a nonpolar solvent):

    • Which dissolves better, H3C-F (Fluoromethane) or H3C-CH3 (Ethane)? Explain.

      • H3C-CH3 (ethane) will dissolve better in benzene. Benzene is nonpolar. Ethane is also nonpolar and will interact with benzene primarily through LDFs, following the "like dissolves like" rule. Fluoromethane is polar and has dipole-dipole forces, which are not as favorable for interacting with a nonpolar solvent like benzene.

Concepts of Gas Pressure

  • Pressure: Defined as Force/Area (P = F/A). In the context of gases, pressure arises from the collisions of gas molecules with the walls of their container. Atmospheric pressure is the force exerted by the column of air above a given area.

  • Conversions:

    • 760 mm Hg = 760 torr = 1 atm

    • SI unit for pressure = Pascal (Pa), where 1 Pa = 1 kg/m·s².

    • Exact conversion: 1 atm = 101.325 kPa (kilopascals).

  • STP (Standard Temperature and Pressure): Defined as 1 atm and 273.15 K (0 °C). These conditions are used for comparing properties of gases.

  • States of matter can be influenced by altering temperature and/or pressure. Increasing temperature generally favors the gas phase, while increasing pressure generally favors the condensed (liquid or solid) phases.

Phase Changes and Properties

  • Normal Boiling Point (Tb): The temperature at which the vapor pressure of a liquid equals 1 atm (standard atmospheric pressure). At this point, the liquid rapidly turns into a gas throughout its bulk.

  • Normal Melting Point (Tm): The temperature at which a solid converts into a liquid at 1 atm (standard atmospheric pressure). At this temperature, the solid and liquid phases coexist in equilibrium.

  • Phase Boundaries: Lines that separate different states of matter on a phase diagram. Points on these boundaries indicate conditions of temperature and pressure where two phases coexist in dynamic equilibrium (the rate of transition from one phase to another is equal to the rate of the reverse transition).

  • Triple Point: The specific conditions (a unique temperature and pressure) where all three phase boundaries meet, allowing all three phases (solid, liquid, and gas) to coexist in dynamic equilibrium.

Critical Properties of Substances

  • Critical Temperature (Tc): The temperature above which a substance cannot exist as a liquid, regardless of how much pressure is applied. Above Tc, the kinetic energy of molecules is too high for IMFs to hold them in a condensed liquid state.

  • Critical Pressure (Pc): The minimum pressure required to condense a gas into a liquid at its critical temperature. Above Tc and Pc, the substance exists as a supercritical fluid.

  • Supercritical Fluid: A state of matter that exists above its critical temperature and pressure. It exhibits properties intermediate between a gas and a liquid: it can diffuse through solids like a gas and dissolve materials like a liquid. Often used as a solvent in green chemistry.

Problems on Phases and Critical Points

(These problems require access to a specific phase diagram for their exact answers. Below are general approaches to solving them assuming a phase diagram for a substance is given.)

  • Determine the phase at specific conditions (e.g., At 0.5 atm and 100 K):

    • Locate the given temperature on the x-axis and pressure on the y-axis of the phase diagram. The region where these coordinates intersect indicates the phase.

  • Find normal boiling and melting points:

    • Normal boiling point is the temperature where the liquid-gas phase boundary crosses the 1 atm pressure line.

    • Normal melting point is the temperature where the solid-liquid phase boundary crosses the 1 atm pressure line.

  • Find the temperature and pressure at the triple point:

    • Locate the single point on the phase diagram where all three phase boundaries (solid-liquid, liquid-gas, solid-gas) converge.

  • Determine the temperature and pressure at the critical point:

    • Locate the endpoint of the liquid-gas phase boundary. This point represents the critical temperature and critical pressure.

  • Determine the phase at 88 atm and 300 °C (example specific values):

    • Convert temperature to Kelvin (300 °C = 573.15 K). Find the intersection on the phase diagram.

  • Find the temperature and pressure at the critical point and triple point (again, requires a specific phase diagram to read the values).