Intermolecular Forces Notes

Intermolecular Forces

• Apply VSEPR theory to determine the shape and bond angles of linear, bent, trigonal planar, tetrahedral, and pyramidal molecules.
• Determine the polarity of molecules using molecular shape, symmetry, and electronegativity.
• Explain the relationship between vapor pressure, melting point, boiling point, and solubility with intermolecular forces (dispersion forces, dipole-dipole attractions, and hydrogen bonding) in molecular covalent substances.

Covalent Bonding (Prior Knowledge)

• Understand that covalent molecules complete their valence electron shell by sharing valence electrons.
• Write the formula of a covalent molecule given the name/vice versa.
• Draw the Lewis dot structure of a molecule for a covalent compound given the formula and/or name.

VSEPR Theory

• Predict the shape of simple molecules by looking at the valence electron pairs around the central atom.

Reminder: Lewis Dot Diagrams

• Complete the Lewis dot diagram quiz at http://www.sciencegeek.net/Chemistry/Review/LewisStructures/
• Understand what a structural formula is.

Bonding and Nonbonding Electrons

• Two types of electron clouds:
  • Bonding electrons: Electrons shared in a covalent bond.
  • Non-bonding electrons (lone pairs): Electrons that belong to the central atom and are not used in bonding.
• Electrons repel one another, so they arrange themselves as far apart as possible.
• Examples:
  • $NH_3$ with 1 lone pair on the central N atom.
  • $H_2O$ with 2 lone pairs on the central O atom.
  • $HCl$ with 3 lone pairs on the central Cl atom.

Predicting the Shape of Simple Molecules

• Molecules form 3D shapes depending on:
  • The number of bonds.
  • The type of bond (single, double, or triple).
  • The number of non-bonding pairs of electrons (lone pairs).
• To determine the shape, identify the central atom first.

VSEPR Theory and Molecular Geometry

• Structures with no lone pairs around the central atom have specific shapes and bond angles.

• Structures with one lone pair around the central atom affect the molecular shape and bond angles.

• Structures with two lone pairs around the central atom affect the molecular shape and bond angles even more.

Lewis Structure for Polyatomic Molecules

• Step 1: Sum the valence electrons from all atoms shown in the molecular formula.
• Step 2: Write atomic symbols for all atoms to show which atoms are connected. The central atom is the most electropositive.
• Step 3: Complete octets for all peripheral atoms bonded to the central atom.
• Step 4: Place any leftover electrons on the central atom.
• Step 5: If there are not enough electrons to give the central atom an octet, form multiple bonds.
• Example: $NF_3$
  • N has 5 valence electrons, F has 7.
  • Sum of valence electrons = $5 + (3 Imes 7) = 26$.
  • N is the central atom.

Lewis Structure for Polyatomic Ions: Boron Tetrafluoride ($BF_4^-$

• Step 1: Sum the valence electrons from all atoms and add or subtract for the charge.
• Step 2: Write atomic symbols showing connections, with the central atom being the most electropositive.
• Step 3: Complete octets for peripheral atoms.
• Step 4: Place leftover electrons on the central atom.
• Remember to put square brackets around the structure and the charge of the ion.
• Example: $BF_4^-$
  • B has 3 valence electrons, F has 7.
  • Sum of valence electrons = $3 + (4 Imes 7) + 1 = 32$ (add 1 for the negative charge).
  • B is the central atom.

Lewis Structures for Polyatomic Molecules: Carbonyl Fluoride ($COF_2$)

• Step 1: Sum valence electrons.
• Step 2: Write atomic symbols, central atom is most electropositive.
• Step 3: Complete octets for peripheral atoms.
• Step 4: Place leftover electrons on the central atom.
• Step 5: If needed, form multiple bonds to give the central atom an octet.
• Example: $COF_2$
  • C has 4 valence electrons, O has 6, F has 7.
  • Sum of valence electrons = $4 + 6 + (2 Imes 7) = 24$.
  • C is the central atom.

3D Molecules Using VSEPR Theory

• 3D shapes depend on the total number of electron clouds around the central atom.
• Two types of electron clouds:
  • Covalent bonds
  • Lone pairs
• VSEPR theory states that electron clouds move as far away from each other as possible to minimize repulsion.
• Use Lewis structures to determine the number and types of electron clouds.

Effect of Repulsion in Molecules

• Greatest repulsion: lone pair - lone pair.
• Then lone pair - bond pair.
• Least repulsion: bond pair - bond pair.

Shapes of Molecules - Linear

• 2 electron clouds, 2 covalent bonds, 0 lone pairs.
• Shape: linear.
• Bond angle: $180^ Imes$.

VSEPR Theory - Linear

• Hydrogen cyanide ($HCN$): H-C≡N, linear, $180^ Imes$.
• Carbon dioxide ($CO_2$): O=C=O, linear, $180^ Imes$.

Shapes of Molecules - Trigonal Planar

• 3 electron clouds, 3 covalent bonds, 0 lone pairs.
• Shape: trigonal planar.
• Bond angle: $120^ Imes$.

VSEPR Theory - Trigonal Planar

• Boron trifluoride ($BF_3$): trigonal planar, $120^ Imes$ bond angles.

Shapes of Molecules - 4 Electron Clouds

• 4 electron clouds, 4 covalent bonds, 0 lone pairs: tetrahedral, $109^ Imes$.
• 4 electron clouds, 3 covalent bonds, 1 lone pair: pyramidal, $107^ Imes$.
• 4 electron clouds, 2 covalent bonds, 2 lone pairs: bent (V-shape), $104.5^ Imes$.

VSEPR Theory - Tetrahedral

• Methane ($CH_4$): tetrahedral, $109.5^ Imes$, bonds are evenly spaced.

VSEPR Theory - Pyramidal

• Ammonia ($NH_3$): trigonal pyramidal, $107^ Imes$, less repulsion between bonding pairs.

VSEPR Theory - Bent

• Water ($H_2O$): bent (V-shape), $104.5^ Imes$, 2 lone pairs at the top of oxygen repel bonded electrons.

Summary Table of Molecular Shapes

• Linear: 2 electron clouds, 2 covalent bonds, 0 lone pairs, $180^ Imes$.
• Trigonal planar: 3 electron clouds, 3 covalent bonds, 0 lone pairs, $120^ Imes$.
• Tetrahedral: 4 electron clouds, 4 covalent bonds, 0 lone pairs, $109^ Imes$.
• Pyramidal: 4 electron clouds, 3 covalent bonds, 1 lone pair, $107^ Imes$.
• Bent: 4 electron clouds, 2 covalent bonds, 2 lone pairs, $104.5^ Imes$.

Examples of Molecular Shapes

• $CH_4$: tetrahedral, 4 bonded electron pairs.
• $NH_3$: trigonal pyramidal, 3 bonded and 1 lone electron pair.
• $H_2O$: bent, 2 bonded and 2 lone electron pairs.
• $CO_2$: linear, bonded electron pairs.
• $BeCl_2$: linear, 2 bonded electron pairs.
• $BF_3$: trigonal planar, 3 bonded electron pairs.

Quiz on Molecular Shapes

• Predict the molecular shape of hydrogen iodide, boron trihydride, nitrogen triiodide, and fluoromethane ($CH_3F$).

Quiz on Ethanoic Acid

• Evaluate the structure of ethanoic acid: number of central atoms, shapes included, and whether it has a single shape.

Learning Check

• Explain the difference between tetrahedral, pyramidal, and bent molecular shapes.
• Describe the difference between single, double, and triple bonds.
• Determine the shapes of various molecules using Lewis structures and perspective diagrams.

Polarity

• Covalent bonds are classified as polar depending on the electronegativity of the atoms involved.
• Refer to the table of electronegativities in the formula and data book.

Bonding and Electronegativity

• Pure (nonpolar) covalent bond: electrons shared equally (electronegativity difference 0.0-0.4).
• Polar covalent bond: electrons shared unequally (electronegativity difference 0.4-2.0).
• Ionic bond: electron transferred (electronegativity difference > 2.0).

Examples of Bond Polarity

• Nonpolar covalent bond: $F_2$ contains a non-polar bond.
• Polar covalent bond: $HF$ contains a polar bond; fluorine is more electronegative.
• Ionic bond: $NaCl$ complete transfer of electrons.

Nonpolar Covalent Bonds

• If the bonded atoms are close in electronegativity (difference of 0.4 or below), both attract the bonding pair of electrons equally.

Polar Covalent Bonds

• If one atom is significantly more electronegative (difference above 0.4), electrons are pulled closer, creating partial charges and a permanent dipole.

Determining Bond Polarity

• If the electronegativity difference (ΔEN) is zero, then the bond is nonpolar covalent.
• If the ΔEN is less than 0.5, the bond is slightly polar.
• If the ΔEN is equal to or greater than 0.5 and 1.6:
  • Ionic if one of the atoms is a metal.
  • Polar covalent if both atoms are non-metals.
• A compound with a metal and a non-metal is considered ionic regardless of ΔEN.

Determine the Polarity of Bonds

• Examples: N-H, O=O, H-Cl, Na-Cl, C-H, S=O, Br-F, F-H, Mg-S.

Molecular Polarity

• For the following molecules: Draw Lewis dot diagram, Determine the molecular structure, Identify if the bonds are polar or nonpolar: N2, HCl, CO2, H2O.

Polarity Examples

• $N_2$: Nonpolar covalent bond.
• $HCl$: Polar covalent bond.
• $CO_2$: Nonpolar.
• $H_2O$: Polar.

Nonpolar Molecules

• A
  molecule is non-polar if it has only non-polar bonds.
• A molecule can also be non-polar if it has polar bonds but the dipoles cancel out due to the symmetry of the molecule.

Polar Molecules

• A molecule is polar if it has polar bonds and is not symmetrical.

What Makes a Molecule Polar?

• A polar molecule has slight positive and negative charges on opposite ends due to asymmetry.
• Determine molecular shape, annotate polarity, and assess symmetry.
• Examples: $PCl<em>3$, $CCl</em>4$, $BH<em>3$, $CH</em>3Cl$, $CH_2O$.

Learning Check on Polarity

• Electronegativity trends on the periodic table.
• Explain the trends in electronegativity.
• Classify bonds in perspective diagrams as nonpolar covalent, polar covalent, or ionic, labeling partial charges.
  • Examples: bromine gas, water, methane, salt, ammonia, fluoromethane ($CH_3F$).

Additional Learning Check

• Identify 2D methods of drawing chemical structures and their advantages/disadvantages.
• Describe advantages/disadvantages of 3D structures.
• Construct Lewis structures and perspective drawings, identify geometry, label bond polarities, and classify molecules as polar or nonpolar.

Intramolecular vs. Intermolecular Forces

• Intra means within the molecule, inter means between the molecules.
• Example: $HCl$ molecules with intramolecular and intermolecular attractions.

Intermolecular Forces

• Three types of intermolecular forces in covalent substances:
  • Dispersion forces (non-polar molecules).
  • Dipole-dipole forces (polar molecules).
  • Hydrogen bonds (some polar molecules).

Intermolecular Forces (IMF)

• As molecules approach each other, their dipoles align, generating attraction. The strength depends on the duration of the dipole.

Dispersion Forces

• Occur between all molecules, polar and non-polar.
• Temporary dipoles form as electrons move around.

Dispersion Forces Explained

• Temporary dipoles induce dipoles in non-polar molecules.
• Weak, brief attraction occurs between these dipoles.

Factors Affecting Dispersion Forces

• More electrons increase the probability of dipole formation.
• Greater surface area allows more contact between molecules.

Dipole-Dipole Attractions

• Occur between molecules with permanent dipoles.
• Stronger than dispersion forces due to permanent dipoles.

Factors Affecting Dipole-Dipole Attractions

• More polarized bonds lead to larger dipole moments and stronger attractions.

Hydrogen Bonds

• Hydrogen is strongly attracted to lone pairs on O, N, or F of another molecule.
• The attraction is stronger than ordinary dipole-dipole interactions.

Polar Molecules Recap

• Permanent dipoles lead to dipole-dipole interactions or hydrogen bonding.
  • Dipole-Dipole Interaction: Occurs between all polar molecules not meeting hydrogen bond conditions.
  • Hydrogen Bonding: Occurs between polar bonds when F/O/N are bonded to H.

Determine Interaction Between Substances

• Examples: $NF<em>3$, $C</em>2H<em>2$, $CF</em>4$, $CH<em>3Cl$, $CH</em>2Cl<em>2$, $H</em>2O$, $HCl$.

Learning Check on Intermolecular Forces

• Describe polarity and distinguish between polar and nonpolar molecules.
• Contrast intramolecular and intermolecular forces.
• Compare dispersion forces, dipole-dipole forces, and hydrogen bonds.
• Identify molecules that have each type of force.
• Predict interactions between substances.

Review of Intermolecular Forces

• Dispersion Forces: Temporary dipoles.
• Dipole-Dipole Attractions: Attractions between positive and negative regions of molecules.
• Hydrogen Bonding: Strongest, between hydrogen and electronegative atoms.

Flowchart for Determining Intermolecular Forces

• Polar molecules can have dipole-dipole forces or hydrogen bonds. Nonpolar molecules have London dispersion forces.

Strength of Forces Holding Substances Together

• Melting and boiling points are determined by intermolecular forces.
  • Network covalent bonds are strongest.
  • Then ionic, metallic, hydrogen bonds, dipole-dipole, and dispersion forces.

Melting and Boiling Points

• Covalent molecules can be symmetrical or asymmetrical. Asymmetrical molecules may participate in hydrogen bonding or dipole dipole interaction. Size dictates melting point.

Melting and Boiling Points Explained

• Melting point is the temperature at which a solid becomes a liquid.
• Boiling point is the temperature at which a liquid becomes a gas.

Factors in Melting and Boiling Points

• Kinetic energy of particles (more energy, more movement).
• Strength of intermolecular forces (greater force, more energy needed to break apart).

Example: Highest Melting Point

• Comparing carbon dioxide and water, consider symmetry, hydrogen bonding, and size.
• Comparing carbon dioxide and methane, consider symmetry, hydrogen bonding, and size.

Analysing Boiling Point Trends

• Boiling points of various hydrides (H2O, HF, NH3, BiH3, H2Te, H2S, SbH3, HI, HCl, HBr, PH3).

Practice Problems on Melting Points

• Ethane, methanol, and fluoromethane in order of increasing melting point.
• Ethane, carbon dioxide, and methane in order of increasing melting point.
• Which solidifies first when HCl and H2O are cooled?
• Melting points of sodium chloride, water, and methane.
• Boiling points of carbon tetrachloride and trichloromethane.

Answers to Practice Problems

• Solutions and explanations for each problem.

States of Matter

• Solid, liquid, and gas phases with transitions like melting, boiling, sublimation, deposition, condensation, and evaporation.

What is Boiling?

• Evaporation vs. Boiling: bubbles can only form in boiling when vapor pressure overcomes atmospheric pressure.

Vapour Pressure

• Vapour pressure is the pressure of a vapour in contact with its solid or liquid form.

Atmospheric Pressure and Boiling

• The water only boils when the vapor pressure becomes equal to the atmospheric pressure.

Vapour Pressure Explained

• At a set temperature, some molecules have enough kinetic energy to transition to a vapour.
• Weaker IMF = more molecules with enough kinetic energy = higher vapour pressure.
• VP means Vapour Pressure, IMF means Intermolecular forces. Low IMF gives High VP/ High IMF gives Low VP
• Temperature measures average kinetic energy.

Vapour Pressure and Intermolecular Forces

• Covalent molecules most likely to evaporate have high kinetic energy and weak intermolecular forces.
• Volatile substances readily vaporize.

Boiling Point (Atmospheric = Vapour)

• High IMF = =lower VP higher boiling temperature/
• Low IMF = higher VP = lower boiling temperature.

Analysing Vapour Pressure Trends

• Vapour pressure curves for various substances (diethyl ether, bromine, ethanol, water, n-octane, ethylene glycol, mercury).

Vapour Pressure Trends Quiz

Determine to what trends are shown in a graph.
Explain Increasing number of carbons increases the boiling point on alkanes.
EXplain Alcohols always have higher boiling points than alkanes.

Solubility

• "Like dissolves like": polar attracts polar, nonpolar attracts nonpolar.
  New attractions formed break attractions and compensate to one another.

Solubility Defined

• Solubility is the ability of a solute to dissolve in a solvent.
• Measured as the maximum amount of solute in a solvent, resulting in a saturated solution.

Solubility and Intermolecular Forces

• Covalent compounds dissolve in solvents with similar properties.
• Polar substances dissolve in polar solvents (miscible).
• Nonpolar substances dissolve in nonpolar solvents (miscible).
• Non-polar and polar substances are immiscible.

Solubility and Intermolecular Forces Example

• Ions dissolve in water molecules, not oil.

Analysing Solubility Trends: Table of different Covalent compounds and their respective Solubility in water at SLC (mL/100 mL).

Learning Check on Solubility

• Explain the relationship between kinetic energy, melting/boiling point, and intermolecular forces.
• Explain the relationship between vapour pressure, boiling point, and intermolecular forces.
• Predict the solubility of methane or ammonia in water.
• Predict whether glucose will dissolve in water or decane based on its molecular structure.

Predicting Trends in Melting and Boiling Points

• Alkanes and their predicted melting and boiling point trends.

Predicting Trends in Vapour Pressure

• Alcohols and their predicted vapour pressure trends.

Sequencing Compounds by Solubility

• Sequence compounds (chlorine, arsine, water, hydrogen sulfide, pentane) from least to most soluble in water and hexane, justifying each answer.

Solubility Quiz

• Substance X is made of crystal waxes that melt readily at 66 degrees. What is one of the compounds that X can be?
• Vapour pressure is greatest when:
  minimal evaporation happens
  atmospheric pressure increases
  a substance has weak intermolecular forces happens
  a substance has low kinetic energy
• Show the molecule structure on a glucose molecule and show if it will dissolves in Either:
  water
  pentane