Exam 2 slideshow based knowt

Kinetic Molecular Theory — overview

The Kinetic Molecular Theory (KMT) explains ideal gas behavior based on a set of assumptions about particle motion and energy; it's the basis for gas laws linking P, V, T, and n.

KMT assumption — particle spacing

Gas particles are far apart; most of the volume of a gas is empty space.

KMT assumption — motion

Gas particles are in constant, random motion with a range of speeds.

KMT assumption — attractive forces

Gas particles have negligible attractive forces between them (ideal gas assumption).

KMT assumption — kinetic energy and temperature

Gas particle kinetic energy is directly proportional to absolute temperature; higher T → higher average speed.

Ideal Gas Law formula

PV = nRT (P = pressure, V = volume, n = moles, R = gas constant, T = kelvin).

Ideal gas constant values (common)

R = 0.082057 L·atm·mol⁻¹·K⁻¹ = 8.3145 J·mol⁻¹·K⁻¹; know which units match P and V.

Combined gas law (n constant)

P₁V₁/T₁ = P₂V₂/T₂ — use when the amount of gas (n) is constant and P, V, T change.

Boyle's law (definition)

a scientific principle stating that for a fixed amount of gas at a constant temperature, the volume of the gas is inversely proportional to the pressure it exerts. At constant T and n, P ∝ 1/V; P₁V₁ = P₂V₂.

Charles's law (definition)

at a constant pressure, the volume of a gas is directly proportional to its absolute temperature, At constant P and n, V ∝ T (in kelvin); V₁/T₁ = V₂/T₂.

Gay-Lussac's law (definition)

states that for a fixed amount of gas at a constant volume, the pressure is directly proportional to its absolute temperature. In simple terms, if you heat a gas in a sealed container, the pressure will increase. At constant V and n, P ∝ T; P₁/T₁ = P₂/T₂.

Units of pressure and conversions

1 atm = 760 mmHg (torr) = 101.325 kPa = 101325 Pa.

Molar volume at STP (useful)

At STP (0°C, 1.00 atm) ideal gases ≈ 22.414 L per mole.

Real gas deviations

Real gases deviate at high pressure and low temperature because particle volume and intermolecular forces become significant.

Gas law problem approach

Always convert temperatures to kelvin, volumes to liters (if using R=0.08206), pressures to matching units; check if n is constant.

Intermolecular forces — definition

Attractive forces between molecules; they determine many physical properties like BP, MP, vapor pressure, viscosity, and solubility.

Types of attractive forces — list

Common types: ionic (ion-ion), ion–dipole, hydrogen bonding, dipole–dipole, London dispersion (induced dipole).

Relative strength of IMFs (general)

General order (strongest → weakest): ionic/ion–ion ≈ ion–dipole > hydrogen bonding > dipole–dipole > London dispersion (depends on size).

London dispersion forces (what & when)

Temporary induced dipoles due to electron motion; present in all molecules but dominant in nonpolar molecules; strength increases with molar mass and surface area.

Dipole–dipole interactions

electrostatic forces of attraction that occur between polar molecules, where the partially positive end of one molecule is attracted to the partially negative end of another, present in polar molecules with significant electronegativity differences

Hydrogen bonding — criteria

H-bonds occur when H is covalently bonded to N, O, or F (donor) and is attracted to a lone pair on N, O, or F (acceptor).

Hydrogen bonding — biological importance

Hydrogen bonding stabilizes DNA base pairing, secondary/tertiary protein structure, and properties of water.

Ion–dipole interactions

attractive forces between an ion and a polar molecule. These interactions are common in solutions of ionic compounds dissolved in polar solvents, like dissolving salt NaCl

Ionic interactions (ion-ion)

the electrostatic attraction between oppositely charged ions, which are formed when one or more electrons are transferred from one atom to another

Effect of IMFs on boiling point

Stronger intermolecular forces → higher boiling point (more energy required to separate molecules into gas).

Effect of IMFs on vapor pressure

Stronger intermolecular forces → lower vapor pressure at a given temperature (molecules less likely to escape to gas).

Effect of IMFs on melting point

Stronger and more numerous attractive forces → higher melting point.

Predicting BP of alkanes trend

For alkanes: longer straight chains → higher BP; increased branching → lower BP (branching reduces surface contact).

Alkane boiling point trend example

Methane −161°C, Ethane −89°C, Propane −42°C, Butane −0.5°C, Pentane 36°C, Hexane 69°C, etc. — BP rises with carbon count.

Surface area and London forces

More surface area → more contact between molecules → stronger London dispersion forces → higher BP/MP.

Melting point trends general

The more symmetric and the stronger the intermolecular attractions, the higher the melting point (also depends on packing).

“Like dissolves like” — meaning

Polar solvents dissolve polar solutes (and ionic solutes); nonpolar solvents dissolve nonpolar solutes. Match polarity and types of attractive forces.

Hydrophilic vs hydrophobic

Hydrophilic: water-loving (polar/ionized, soluble in water). Hydrophobic: water-fearing (nonpolar, insoluble in water).

Solubility of ionic compounds in water

Water hydrates ions: multiple ion–dipole interactions with water overcome ionic lattice energy → many ionic salts are soluble.

Solubility of nonpolar oils in water

Nonpolar oils interact only via London forces; water's hydrogen bonding/dipole–dipole interactions are stronger among water molecules → oil and water are immiscible.

Sucrose solubility in water

Contains many hydroxyl groups; can hydrogen-bond with water → soluble.

Amphipathic molecules definition

Molecules with both polar (hydrophilic) and nonpolar (hydrophobic) regions (e.g., fatty acids, phospholipids).

Fatty acid polarity

Fatty acids have a polar carboxyl head and a long nonpolar hydrocarbon tail; large tail often dominates making them mostly nonpolar.

Soaps and micelles

Soaps are fatty acid salts (ionic head, nonpolar tail). In water they form micelles: hydrophobic tails inward, ionic heads outward interacting with water.

Why micelles solubilize oils

In a micelle, hydrophobic tails solubilize nonpolar solutes (oil) in the core while ionic heads interact with water — enables oil removal.

Pharmaceutical solubility strategy

Convert poorly water-soluble compounds to charged forms (e.g., protonate an amine with HCl) to increase ion–dipole hydration and water solubility.

Pseudoephedrine → salt example

Reacting pseudoephedrine (base) with HCl produces a protonated salt that hydrates and dissolves more readily in water.

Phospholipid bilayer — structure

Phospholipids form bilayers with polar heads facing aqueous environments (inside and outside cell) and hydrophobic tails inward forming a nonpolar core.

Cholesterol’s role in membranes

Cholesterol intercalates in bilayers, decreasing membrane fluidity at high T and preventing solidification at low T; reduces permeability.

Cholesterol distribution

Cholesterol is present in animal cell membranes; plant cells do not contain cholesterol (they have other sterols).

Chapter 7 quick study tip

CH7 focuses on conceptual understanding of IMFs, solubility trends, cell membranes, and gas-law relationships — expect conceptual multiple-choice.

Solution definition (chemistry)

A solution = solute(s) dissolved in a solvent; an aqueous solution has water as the solvent.

Solvent vs solute definitions

Solvent = component present in greatest amount; solute = component(s) present in lesser amounts.

Aqueous solution

solution where water is the solvent

Solubility vs saturation

Solubility = maximum amount of solute that dissolves in a given amount of solvent at a specified temperature; saturated = holds the maximum amount.

Temperature effect — solids in liquids

For most solids, solubility increases with temperature (you can dissolve more sugar in hot tea).

Temperature effect — gases in liquids

Solubility of gases in liquids decreases as temperature increases (CO₂ leaves warm soda).

Henry's law statement

The solubility of a gas in a liquid is directly proportional to the partial pressure of that gas above the liquid: S = kH × Pgas.

Henry's law application (soda)

CO₂ solubility in soda increases with CO₂ pressure; opening the bottle lowers P_CO2 → CO₂ comes out of solution as bubbles.

Electrolyte definitions

Strong electrolytes fully dissociate into ions (conduct electricity well); weak electrolytes partially dissociate; nonelectrolytes do not produce ions in solution.

Examples of strong electrolytes

Strong acids (HCl, HBr, HNO3), strong bases (NaOH, KOH), and soluble ionic salts (NaCl, KBr).

Nonelectrolyte

a substance that does not form ions when dissolved in water, and therefore, its solution does not conduct electricity EX: Sucrose

Hydration (definition)

When ions are surrounded by water molecules, ion–dipole attractions stabilize them in solution — called hydration (or solvation).

Concentration units — molarity

M = moles solute per liter of solution (mol/L).

Concentration units — mass percent

% (w/w) = mass solute / mass solution × 100%.

Concentration units — mass/volume percent

% (m/v) = mass solute (g) per volume solution (mL) × 100% (common in biology/med).

Concentration units — volume percent

% (v/v) = volume solute / volume solution × 100% (used for liquid mixtures).

mM and mmol/L

mM = millimolar = mmol/L (useful for electrolytes and clinical concentrations).

Equivalents and mEq

An equivalent = moles × charge; mEq/L = (mmol/L) × |charge|. For Ca²⁺, 1 mmol = 2 mEq.

Molarity example — method

Given mass% and density, assume 1.00 L solution → compute mass of solution, mass of solute (mass% × total mass), moles solute, then moles ÷ volume(L) = M.

Worked molarity example — 30.0% (w/w) H₂O₂ with D=1.11 g/mL

Assume 1.00 L solution = 1000 mL; mass = 1.11 g/mL × 1000 mL = 1110 g; H₂O₂ mass = 0.300 × 1110 g = 333 g; moles = 333 g ÷ 34.0138 g/mol ≈ 9.79 mol; M ≈ 9.79 M.

Dilution equation

M₁V₁ = M₂V₂ — used to dilute a concentrated stock to a desired concentration; volume units must match.

Dilution example — prepare 1.00 L of 0.10 M from 1.00 M stock

V₁ = (M₂V₂)/M₁ = (0.10 M × 1.00 L) / 1.00 M = 0.10 L = 100 mL stock; add water to 1.00 L.

Mixing units trick

For molarity problems, pick convenient sample size (1.00 L or 1000 g) to make percent conversions easy.

Diffusion definition

Movement of molecules from high concentration to low concentration until equilibrium; driven by concentration gradient.

Osmosis definition

Net movement of water across a semipermeable membrane from region of higher water concentration (dilute) to lower water concentration (concentrated).

Isotonic/hypotonic/hypertonic definitions

Isotonic: equal solute concentration → no net water movement; Hypotonic: lower solute outside → water enters cell (may lyse); Hypertonic: higher solute outside → water exits cell (cell shrinks).

Clinical isotonic solution

RBCs: 0.9% (m/v) NaCl is isotonic for red blood cells; IVs often use this or 5% dextrose depending on clinical need.

Membrane transport types — simple diffusion

Small nonpolar molecules (O₂, CO₂, N₂, some small lipids) diffuse directly through the membrane down their concentration gradient.

Membrane transport types — facilitated diffusion

Carrier or channel proteins assist polar/charged molecules (e.g., glucose transporters); still down gradient and does not require ATP.

Membrane transport types — active transport

Transport against a concentration gradient using energy (usually ATP) via pumps (e.g., Na⁺/K⁺ ATPase).

Chapter 8 quick study tip

Know how temperature and pressure affect solubility, common concentration units and conversions, dilution math, and membrane transport types.

Arrhenius acid/base definitions

Acid: increases [H⁺] in water; Base: increases [OH⁻] in water (Arrhenius).

Bronsted–Lowry acid/base definitions

Acid: proton (H⁺) donor; Base: proton acceptor. Conjugate pairs are acid/base pair differing by one proton.

Conjugate acid–base pair example

HCl (acid) ↔ Cl⁻ (conjugate base); NH₃ (base) ↔ NH₄⁺ (conjugate acid).

Strong acids — common list

Common strong acids: HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄ (first proton strong).

Strong bases — common list

Common strong bases: NaOH, KOH, LiOH, Ca(OH)₂ (sparingly soluble in water but strong), Sr(OH)₂, Ba(OH)₂.

pH definition

pH = −log₁₀[H⁺]; measures acidity; lower pH = more H⁺.

pOH definition

pOH = −log₁₀[OH⁻]; pH + pOH = 14.00 at 25°C (approx).

Calculating [H⁺] from pH

[H⁺] = 10^(−pH).

Henderson–Hasselbalch equation

pH = pKa + log₁₀([A⁻]/[HA]) — useful for buffer pH calculations and buffer design.

pKa meaning

The negative log of the acid dissociation constant (Ka); smaller pKa = stronger acid.

Weak acid dissociation

the partial, reversible ionization of a weak acid in water, where only a small fraction of the acid molecules break apart into ions and reach an equilibrium state

Neutralization reaction

A neutralization reaction is when an acid and a base react to form water and a salt and involves the combination of H + ions and OH - ions to generate water.

Chemical equilibrium definition

At equilibrium forward and reverse rates are equal; concentrations remain constant though both species are present.

Equilibrium constant K

a value that describes a chemical reaction at equilibrium, representing the ratio of products to reactants, each raised to the power of its coefficient in the balanced chemical equation

Interpreting K values

K = 1 → products ≈ reactants at equilibrium; K > 1 → products predominate; K < 1 → reactants predominate.

Le Châtelier’s principle

Le Chatelier's principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change to reestablish an equilibrium.

Le Châtelier — adding reactant effect

Add reactant → reaction shifts to products (right) to reduce stress.

Le Châtelier — removing product effect

Remove product → reaction shifts to products (right) to replace lost product.

Le Châtelier — pressure change (gases)

Increasing pressure (by decreasing volume) shifts equilibrium toward the side with fewer moles of gas; decreasing pressure → side with more gas moles favored.

Le Châtelier — adding inert gas (constant V)

Adding inert gas at constant volume does not change partial pressures of reactants/products; no shift in equilibrium.

Le Châtelier — temperature change & K

Temperature changes change K: raising T favors endothermic direction (K changes accordingly); K is only constant at a fixed temperature.

Bicarbonate buffer system — chemical sequence

CO₂(g) + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻; this system resists pH change in blood.

Henderson–Hasselbalch for bicarbonate

pH = pKa + log([HCO₃⁻]/[H₂CO₃]) — clinically H₂CO₃ is often related to pCO₂.