BCHEM 153 Lecture Notes Flashcards

BCHEM 153: General Chemistry Notes

Class Plan & Reminders

• Class Plan:
  • Check-in and reminders (5 min)
  • Review of Intermolecular Forces (IMFs) (15 min)
  • Intermolecular Forces/Boiling Point discussion (45 min)
  • Wrap-up (10 min)
• Materials:
  • Sit in POGIL teams
  • Pencil
  • Calculator
  • Electronic device
• Reminders:
  • ALEKS Chapter 10 is now open.
  • Quiz this week is a mid-quarter assessment.
  • Lab this week is Heat of Reaction - lab reports are due.
  • Meet with peer facilitator.
• Office Hours:
  • Dr. Lovitt: MW 12:30-1:45pm in UW2-030 (QSC), by appointment on calendly.com/lovittc
  • Simran: Tuesday 8:15pm to 9:15pm
  • Aleks: 11:00-12:00 on Friday in QSC
  • Willow: Wednesdays after class in QSC
  • Bryson: Tuesday 5:45-6:45pm in QSC
  • Rimjhim: Thursday 3-4 PM on Zoom

Introduction to Intermolecular Forces and Boiling Points

Learning Outcomes

• Identify intramolecular and intermolecular forces.
• Determine the relative strength of intermolecular forces.
• Relate boiling point trends to intermolecular forces.
• Process Skill: Critical Thinking - Use molecular structures to describe trends in molecular properties.

Intramolecular vs. 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.
• Intermolecular forces have considerable influence on the physical properties of all substances.

Water Molecules and Intermolecular Forces

• Water molecules stick to one another due to attractive intermolecular forces.
• This explains:
  • Why ice is solid
  • Liquid water
  • Why water “beads” on your car
  • Surface tension
  • Capillary action, etc.
• Gaseous water molecules lack these attachments.

Review: Different Types of Bonds

• Pure Covalent: equal sharing of electrons (e.g., $Cl_2$)
• Polar Covalent: unequal sharing of electrons (e.g., $HCl$)
• Ionic: transfer of electrons (e.g., $NaCl$)
• Electronegativity difference (ΔEN) and Bond Character:
  • $ΔEN > 1.7$: Mostly ionic
  • $ΔEN = 0.4-1.7$: Polar covalent
  • $ΔEN < 0.4$: Mostly covalent
  • $ΔEN = 0$: Nonpolar covalent

Review: Polar Bonds and Polar Molecules

• Bond dipole: Separation of electrical charge created when atoms with different electronegativities (EN) form a covalent bond.
• Polar molecule: Vectors of bond dipoles whose sum is greater than zero.

Molecules Explored

• Butane ($C<em>4H</em>{10}$) - Alkane (hydrocarbon containing only C and H with single bonds)
• 1-propanol ($C<em>3H</em>8O$) - Alcohol (contains an O-H group)
• Acetone ($C<em>3H</em>6O$) - Ketone (has a carbon double bonded to oxygen, C=O)

Electronegativity Trends

• Trends:
  • Electronegativity generally increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

Interactions of Nonpolar Molecules: Dispersion Forces

• Dispersion (London) forces: Intermolecular force between nonpolar molecules caused by temporary dipoles.
• Temporary (induced) dipole: Separation of charge produced in an atom or molecule by a momentary uneven distribution of electrons.
• Polarizability: Relative ease with which the electron cloud in a molecule, ion, or atom can be distorted, inducing a temporary dipole.

Boiling Points and Atomic Number

• The boiling point increases as the atomic number increases.
• Separating liquid-phase particles that are attracted to each other requires energy.
• The stronger the particles’ attractions, the greater the energy needed to separate them.
• The greater the energy required to separate these particles, the higher the boiling points.

Induced Dipole Illustration

• Momentary shift in electron density = dispersion
• Temporary dipoles
• Electron-poor vs. electron-rich regions

Induced Dipoles: Sloshy Bathtub Analogy

• No polarization in isolated atoms/molecules.
• Instantaneous dipole on one atom/molecule induces a dipole on another.

Strength of Dispersion Forces

• Factors affecting strength of dispersion:
  1. Size of atoms/molecules:
     • Larger atoms/molecules are more polarizable than smaller ones.
     • Dispersion increases with polarizability.
  2. Shape of molecules:
     • Increased surface area = increased interactions between molecules.
     • Linear molecules have higher dispersion than branched molecules of similar molecular weight.

Effect of Size on Dispersion

• Table showing boiling points of halogens (F2, Cl2, Br2, I2, At2) vs. molar mass.
• Boiling point increases with molar mass due to increased dispersion forces.

Effect of Shape on Dispersion

• The more spread out a molecule, the stronger the opportunity for dispersion forces and the higher the boiling point.

Importance of London Dispersion Forces

• Increases greatly as size of atom or molecule increases.
• Greater polarizability.

Dipole Interactions

• Presence of a dipole increases the amount of attractive forces between molecules.
• Comparison of Butane (C4H10) and Acetone (C3H6O) - Acetone has a higher boiling point due to dipole-dipole interactions.

Dipole-Dipole Interactions and Hydrogen Bonding

• Dipole–dipole interaction: Attractive force between polar molecules.
• Hydrogen bond: Strongest dipole–dipole interaction. Occurs between H atom bonded to a small, highly electronegative element (F, O, N) and an atom of F, O, or N in another molecule.

Attractive Forces Between Water Molecules

• Attractive forces between water molecules are stronger in the liquid than the gas because the molecules are very close to one another.

Boiling Points of Binary Hydrides

• Boiling points are related to attractive forces.
• Increased boiling points of $H<em>2O$, $HF$, and $NH</em>3$ are due to hydrogen bonding.

Hydrogen Bonding

• Special dipole-dipole interactions that happens between O, N, or F and H bonded to O, N, F.

Predicting Boiling Points Practice

• Rank the following compounds in order of increasing boiling point: $CH<em>3OH$, $CH</em>3CH<em>2CH</em>2CH<em>3$, and $CH</em>3CH<em>2OCH</em>3$
• Lowest = $CH<em>3CH</em>2CH<em>2CH</em>3$ (molar mass = 58 g/mol). The only available force is dispersion, which is the weakest force.
• Middle = $CH<em>3CH</em>2OCH_3$ (molar mass = 60 g/mol). Although of similar molar mass as the first compound and having similar dispersion forces, this molecule also contains O and will have a dipole.
• Highest = $CH_3OH$ (molar mass = 32 g/mol). Although this is the smallest molecule, it is capable of H bonding, which is the strongest of the intermolecular forces. We would expect this substance to have the highest boiling point.

Relative Strengths of Intermolecular Forces

• ion-dipole
• hydrogen bonding
• dipole-dipole
• dipole-induced dipole
• dispersion

Hydrogen Bonding in DNA

• H-bonding between complementary strands of DNA.
• Adenine (A) and thymine (T) form two hydrogen bonds.
• Guanine (G) and cytosine (C) form three hydrogen bonds.

Ion-Ion Forces

• Ion-ion forces (only in ionic compounds)
• In a pure ionic compound, ion-ion forces are indistinguishable from ionic bonds.

Interactions Involving Polar Molecules: Ion-Dipole Interaction

• Ion–dipole interaction: Attractive force between an ion and a molecule that has a permanent dipole.

Sphere of Hydration

• Sphere of hydration: Cluster of water molecules surrounding an ion in aqueous solution.
• Sphere of solvation: If solvent is other than $H_2O$

Dipole-Induced Dipole

• Proximity of polar molecule causes dipole-induced dipole

States of Matter

• Solid, Liquid, Gas

Fracking

• Hydraulic fracturing, or "fracking," involves the injection of more than a million gallons of water, sand and chemicals at high pressure down and across into horizontally drilled wells as far as 10,000 feet below the surface. The pressurized mixture causes the rock layer, in this case the Marcellus Shale, to crack. These fissures are held open by the sand particles so that natural gas from the shale can flow up the well.

Interactions in Molecules - Example

• Examples identifying Ion-dipole attraction, Dipole-dipole attraction, Hydrogen bonding, London Dispersion forces between molecules

Substances in Fracking Fluid

• Examples of substances possibly present in fracking fluid (Ethylene glycol, Citric acid, Hydrochloric acid, Potassium chloride, Formic acid, Boric acid, Sodium carbonate)

Hydrogen Bonds in Polyacrylamide Molecule

• Illustration showing hydrogen bonds in polyacrylamide

Solubility

• Solubility depends on relative strength of solute–solvent interactions compared to solute–solute or solvent–solvent.
• Like dissolves like!

Predicting Solubility

• Methanal and butanol have more similar interactions so they are more likely to mix with each other.

Solubility in Water Practice

• Considering the compounds carbon tetrachloride ($CCl<em>4$), ammonia ($NH</em>3$), hydrogen fluoride ($HF$), and oxygen ($O_2$), which should be very soluble in water and which should have limited solubility in water?
• The polar molecules $NH_3$ and $HF$ are soluble in water.
• The nonpolar molecules $CCl<em>4$ and $O</em>2$ have only limited solubility in water.

Solubility Terms

• Miscible
• Soluble
• Hydrophilic
• Immiscible
• Insoluble
• Hydrophobic

Properties of Water

• Surface tension: Energy needed to separate the molecules at the surface of a liquid
• Viscosity: Measure of resistance to flow of a liquid
• Related to intermolecular forces:
  • As interactive forces increase, surface tension and viscosity also increase.
  • Related examples: meniscus, capillary action

Vapor Pressure

• The pressure exerted by molecules of a substance in the gas phase in equilibrium with its liquid phase
• Factors that affect vapor pressure
  1. Temperature: Higher temperature = greater kinetic energy
  2. Surface area: Number of molecules in position to enter gas phase
  3. Intermolecular forces: Greater force requires greater kinetic energy needed for molecule to escape surface

Vapor Pressure Illustration

• Illustration (a) and (b)

Vapor Pressure vs. Temperature

• Examples (Diethyl ether, Ethanol, Water, Ethylene Glycol)

Vapor pressure is related to the temperature-ability of molecules to 'escape liquid phase'

• Lower temperature vs Higher temperature

Surface Tension

• Intermolecular forces holding molecules together

Cohesive vs. Adhesive Forces: Meniscus

• Curvature of liquid surface due to adhesive and cohesive forces
• Concave: adhesive forces ≥ cohesive forces (e.g., water on glass)
• Convex: cohesive forces > adhesive forces (e.g., mercury on glass)

Cohesive vs. Adhesive Forces: Capillary Action

• The rise of a liquid up a narrow tube due to adhesive forces between the liquid and the tube and cohesive forces within the liquid

Viscosity

• Resistance to flow
• Factors:
  • Intermolecular forces
  • Molar mass
  • Molecular shape
  • Temperature

Density

• Density of water decreases as it freezes.
• H–bonding results in a cage-like structure in solid state; less dense than liquid state.

Water and Aquatic Life

• Lakes/rivers freeze from the top down, allowing fish and aquatic life to survive below.
• As surface waters warm or cool, nutrient-rich bottom waters cycle to the surface; oxygen-rich surface waters cycle to the bottom.

What determines state of matter?

• Kinetic energy of molecules
• Intermolecular forces
• The BALANCE of the two determines the state of a given substance at a given temperature

Review: States of Matter

• Physical state of a substance depends on
  • Kinetic energy of particles
  • Intermolecular forces
  • Pressure exerted (forces particles together)

Physical State and Phase Transformations

• Phase diagram:
  • A graphical representation of how the stabilities of the physical states of a substance depend on temperature and pressure
  • Equilibrium lines represent phase changes between solid/liquid, liquid/gas, and solid/gas under specific conditions (T, P).
  • Supercritical region: fourth state of matter with properties intermediate between gas and liquid

Phase Diagram for Water

• Triple point: Temperature/pressure where all three phases of a substance coexist
• Critical point: Specific temperature/pressure at which the liquid and gas phases have the same density
• Supercritical fluid: A substance above its critical temperature and pressure

Phase Diagram for $CO_2$

Phase Diagrams Terms

• Normal: Occurs at 1 atm.
• Critical Point: A combination of temperature and pressure beyond which a gas cannot be liquefied.
• Critical Temperature: The temperature beyond which a gas cannot be liquefied regardless of the pressure.
• Critical Pressure: The pressure beyond which a liquid cannot be vaporized regardless of the temperature.
• Supercritical Fluid: A state of matter beyond the critical point that is neither liquid nor gas.

Phase Diagrams

• Phase diagrams show how pressure and temperature (KE) determine state of the substance

Phase diagram for water

• The solid/liquid boundary line has a negative slope.