Intermolecular Forces: Types and Interactions

Intermolecular Forces Overview

• Focus is on intermolecular forces, including ion-ion, ion-dipole, and dipole-dipole interactions
• Discussion of hydrogen bonds, London dispersion forces (LDF), and van der Waals forces
• Examples provided to identify types of interactions in different compounds

Ion-Ion Interactions

• Occur between oppositely charged ions, like sodium (Na+) and chloride (Cl-)
• Electrostatic force is proportional to the charge and inversely related to the distance
• Higher charges lead to greater ionic interactions (e.g., Calcium (Ca+2) and oxide (O-2) interaction is stronger than Na+ and Cl-)
• Lattice energy is proportional to the magnitude of charges and inversely related to the distance between ions

Lattice Energy and Melting Point

• Compounds with higher lattice energy have higher melting points
• Aluminum nitride (AlN) has higher melting point than magnesium oxide (MgO) because Al+3 and N-3 charges are greater than Mg+2 and O-2
• Sodium fluoride (NaF) has a higher melting point than potassium chloride (KCl) because Na+ and F- are smaller than K+ and Cl-

Ion-Dipole Interactions

• Ions are particles with unequal numbers of protons and electrons
• Atoms have an equal number of protons and electrons; ions have a net charge
• Dipole is a neutral molecule with partial positive and negative charges

Dipole Definition

• Dipoles have two poles of charge; one side is positive, and the other is negative
• Carbon monoxide (CO) is a dipole; oxygen (O) is partially negative, carbon (C) is partially positive
• Water (H2O) is a permanent dipole because oxygen is more electronegative than hydrogen

Ion-Dipole Examples

• Interaction between sodium cation (Na+) and the oxygen atom of water (H2O)
• Interaction between chloride anion (Cl-) and the hydrogen atom of water (H2O)
• Water molecules surround and solvate ions (e.g., Na+ and Cl-) when sodium chloride (NaCl) dissolves in water

Dipole-Dipole Interactions

• Occur between two polar molecules
• Hydrogen atom of one molecule is attracted to the bromine atom of another molecule
• Oxygen atom of one carbon monoxide molecule is attracted to the carbon atom of another

Hydrogen Bonds

• Occur between hydrogen (H) and nitrogen (N), oxygen (O), or fluorine (F)
• Intermolecular interaction between two separate water molecules
• Covalent bonds (O-H) are intramolecular, existing within a single water molecule
• Hydrogen bonds are strong dipole-dipole interactions

London Dispersion Forces (LDF)

• Found in all molecules, most significant in non-polar molecules
• Result from weak, temporary dipole interactions
• Temporary dipoles occur when electron clouds distort, creating momentary charges
• Induced dipoles are created by the presence of other temporary dipoles
• Weaker than typical dipole-dipole interactions

Strength of Intermolecular Forces

• Strongest to weakest: ion-ion > ion-dipole > hydrogen bond > dipole-dipole > LDF
• LDF also known as van der Waals forces, are associated with temporary induced dipoles

Compound Interaction Examples

• Magnesium oxide (MgO): ion-ion interaction
• Potassium chloride in water (KCl in H2O): ion-dipole interaction
• Methane (CH4): London dispersion forces (LDF) due to being non-polar

Additional Compound Examples

• Carbon dioxide (CO2): LDF because it is non-polar due to canceling dipole moments
• Sulfur dioxide (SO2): dipole-dipole due to bent shape and non-canceling dipole moments
• Hydrochloric acid (HF): hydrogen bonds between molecules

More Compound Examples

• Methanol (CH3OH) and lithium chloride (LiCl): ion-dipole interaction
• Formaldehyde (CH2O) and carbon monoxide (CO): dipole-dipole interaction

Boiling Point Comparison: I2 vs Br2

• Both are non-polar, so LDF is the strongest intermolecular force
• Iodine (I2) has more electrons, higher polarizability, and stronger LDF, resulting in a higher boiling point

Ranking Boiling Points: Cl2, I2, Br2, F2

• All non-polar, so LDF is the predominant force
• Boiling point increases with size: F2 < Cl2 < Br2 < I2

Boiling Point: CH3OH vs CH4

• Methanol (CH3OH) is polar with hydrogen bonds; methane (CH4) is non-polar with LDF
• Methanol has higher boiling point due to stronger intermolecular forces

Boiling Point: Propanol vs. Methanol

• Both have hydrogen bonds
• Propanol (CH3CH2CH2OH) is larger, leading to more LDF and a higher boiling point

Vapor Pressure and Volatility: Propanol vs. Methanol

• Methanol has lower boiling point/higher vapor pressure, making it more volatile
• Propanol has higher boiling point/lower vapor pressure

Solubility in Water: Methanol vs. Propanol

• Polar substances dissolve in water
• Methanol is more soluble due to its small non-polar region
• Propanol has a larger non-polar region, reducing its solubility

Solubility and Hydrocarbon Chain Length

• Longer hydrocarbon chains reduce solubility in water
• Octanol (eight carbons) is nearly insoluble
• Methanol and propanol are more soluble because they are relatively polar

Boiling Point: Neopentane vs. Pentane

• Isomers with same molecular weight but different structures
• Straight-chain alkanes (pentane) have higher boiling points
• Branched alkanes (neopentane) have lower boiling points due to smaller surface area

Intermolecular Interactions and Surface Area

• Larger surface area means more contact and stronger temporary induced dipole interactions
• Neopentane has less intermolecular interactions, and a lower boiling point

H2O, H2S, and H2Se Boiling Points Rank

• H2O has hydrogen bonds, making it highest
• H2Se is larger and more polarizable, with a higher boiling point than H2S

Ranking Order of Boiling Point (HF, HBr, HI, HCl)

• Higher Boiling Point: HF > HI > HBr > HCL
• HI largest between HCL and HBr so bowling point is higher
• Hydrogen bonds have the highest bowling point