Chapter 10-17.2 for Midterm

just merged all of my AP Chemistry notes from past tests and quizzes. can find specific chapters in my AP Chemistry folder

Boyle’s Law

the volume of a fixed quantity of gas is inversely proportional to its pressure; V = constant x 1/P or PV = constant

Charles Law

the volume of a fixed quantity of gas at constant pressure increases as the temperature increases; V = constant x T or V/T = constant

Gay-Lussac’s Law of combining volumes

at a given temperature and pressure, the volumes of gases which react are ratios of small whole numbers; the pressure and Kelvin temperature of a gas are directly proportional, provided the volume remains constant; P₁/T₁ = P₂/T₂

Combined Gas Law

relationship between pressure, volume, and temperature of a fixed amount of gas; combination of Boyle’s, Charles's, and Gay-Lussac’s laws; (P₁V₁)/T₁ = (P₂V₂)/T₂

Avogadro’s Hypothesis

equal volumes of gas at the same temperature and pressure will contain the same number of molecules

Avogadro’s Law

the volume of gas at a given temperature and pressure is directly proportional to the number of moles of gas; 22.4 L of any gas at 0 degrees C contain 6.02 × 10²³ gas molecules; V = constant x n

Gases

composed of nonmetallic elements, simple moleculars formulas, and low molar masses; made up of molecules or atoms that are arranged without structure; no fixed shape or volume

Vapor

substances that are liquids of solids under ordinary conditions that also exist in the gaseous state; H₂O exists as liquid water, solid ice, or water vapor

Iiquid

has a definite volume but no definite shape; made up of atoms or molecules that are connected by bonds and their particles can flow freely; less rigid than solids but more rigid than gases

Solid

a substance that has definite shape and volume; particles are arranged in a specific arrangement; firm or hard

Properties of Gas

expand spontaneously (to fill a container); highly compressible; form homogeneous mixtures; nonmetallic elements; different chemical properties

Absolute Zero

0 K or -273.15 degrees C; William Thomson proposed an absolute-temperature scale known as Kelvin scale

Density as it relates to gases

d = (nM)/V = (PM)/(RT); depends on its pressure, molar mass, and temperature; the higher the molar mass and pressure, denser the gas; higher temperature, less dense the gas; less dense gas will lie above a denser gas without mixing (hotter gas is less dense so it rises); molar mass of a gas = M = (dRT)/P

Kinetic Molecular Theory

the theory of moving molecules; pressure of a gas caused by collisions with container but its magnitude is determined by how often and how forcefully the collisions happend; proportional to temperature so at any temperature, same average kinetic energy; if absolute temperature is doubled, average kinetic energy of its molecules doubles; KE = ½ mu²

5 Parts of Kinetic Molecular Theory

random motion (gases consists of molecules in continuous random motion); negligible molecular volume (combined volume of all of the gas’ molecules is relative to its total volume); negligible forces (attractive and repulsive forces between gas molecules are insignificant); constant average kinetic energy(as temperature stays constant, the average kinetic energy of molecules does not change; energy can be transfered during collisions though); average kinetic energy proportional to temperature (proportional to absolute temperature; at any temperature, all gas molecules have same average kinetic energy)

Diffusion

spread of one substance throughout a space for a second substance; faster for light gas molecules; significantly slower than RMS speed; slowed by gas molecules colliding with each other

Effusion

escape of gas molecules through a tiny hole; Graham’s Law

Graham’s Law of Effusion

r₁/r₂ = sqrt(M₂/M₁); r₁ and r₂ - effusion of two gases; M₁ and M₂ - molar masses; a lighter gas has the higher effusion rate; rate of effusion is proportional to the rms speed (root mean square speed- speed of molecule having the same kinetic energy and average kinetic energy)

Maxwell Boltzmann

the distribution of speeds for a gas at a certain temperature; higher temperatures increase average particle speed; high molar mass, average particle speed decreases; average kinetic energy increases as temperature increases because temperature increase increases particle speed that is directly proportional to kinetic energy; molar mass does not affect average kinetic energy because increasing molar mass decreases particle speed which cancel out so kinetic energy stays the same

What causes deviations from the ideal gas law?

higher pressure = greater deviation → gas molecules get closer, intermolecular distance decreases so attractive forces take over; temperature increase = more ideal gas → because gas molecules move faster and apart so more energy is available to break intermolecular forces; increase with increasing molecular complexity (volume + attractive forces) and increasing mass (volume); volumes of real gases are larger and have smaller pressures of ideal gases

Ideal Gas

V = (nRT)/P; molecules do not interact with each other; molecules’ combined volume is much smaller than the volume the gas occupies

Standard Temperature and Pressure

0 C and 1 atm

Pressure (P)

force per unit; SI is pascals (Pa); bars (bar = 10⁵ Pa); atmosphere (atm) and torr (torr or mmHg); barometer measures atmospheric pressure; manometer measures pressure of enclosed gases

Volume (V)

liters

Temperature

kelvins

Dalton’s Law of partial pressure

each gas exerts if present alone under same conditions; add all of the partial pressures up to make total pressure; partial pressure = mole fractions times total pressure

Mole Fraction

ratio of moles of one component of a mixture to the total moles of all components

root-mean-square (RMS) speed, u_rms

varies in proportion to the square root for the absolute temperature and inversely with the square root of the molar mass; = sqrt((3RT)/M) but most probably speed of a gas molecules is u_mp = sqrt((2RT)/M)

Mean free path

mean distance traveled between collisions; moving molecules has short path; collisions between molecules limit diffusion rate

as pressure increases, volume

decreases

as volume increase, pressure

decreases

as temperature increases, volume

increases

as pressure increases, n

increases

Particle arrangement in solids

closely packed in an ordered array; positions are essentially fixed; energies of particle-particle attraction are greater than kinetic energies of particles

Particle arrangement in liquids

particles are closely packed but randomly orientated; retain freedom of motion; kinetic energies of particles similar to energies of particle-particle attraction

Particle arrangment in gas

particles are far apart; posses complete freedom of motion; kinetic energies of particles are greater than the energies of particle-particle attraction

Ion-Dipole Forces

interaction between an ion and a dipole; a neutral, polar molecule (ex: water); strongest of all intermolecular forces (solutions ONLY)

Dipole-Dipole Forces

between neutral polar molecules (oppositely charged ends of molecules attract); weaker than ion-dipole forces; increase with increasing polarity; strength of attractive forces is inversely related to molecular volume

London Dispersion Forces

weakest of all intermolecular forces; two adjacent neutral, nonpolar molecules; the nucleus of one attracts the electrons of the other; electron clouds are distorted; instantaneous dipole; strength of forces is directly related to molecular weight; exist between all molecules; depend on the shape of the molecules; the greater the surface area available for contact, the greater the force is

Hydrogen Bonding

special case of dipole-dipole forces; H-bonding requires H bonded to an electronegative element (F, O, N); boiling increases with increasing molecular weight (exception water)

Intermolecular Forces

London dispersion forces, dipole-dipole forces, hydrogen bonding, ion-dipole forces, ionic bonding

What is the effect of molar mass on IMFs?

increasing molar mass has stronger IMFs

What is the effect of structure on IMFs?

longer molecules have a greater surface area result in stronger IMFS

The stronger the attractive forces, the (boiling point and melting point)

the higher the boiling point of the liquid and the melting point of a solid

Temperature of boiling point increases as pressure

decreases

high altitude: low pressure so water boils at

lower temperature

liquids boil when the external pressure equals

the vapor pressure

Normal boiling point

Boiling point of a liquid at 1 atm

Vapor Pressure

pressure exerted when the liquid and vapor are in dynamic equilibrium; some molecules on the surface of a liquid have enough energy to escape to the gas phase, after some time the pressure of the gas will be constant at the vapor pressure (equilibrium)

Dynamic Equilibrium

the point when as many molecules escape the surface as strike the surface

Vapor Pressure increases nonlinearly with increasing

temperature (Clausius-Clapeyron Equation)

If equilibrium is never established then the liquid

evaporates; volatile substances (high VP) evaporate rapidly

the higher the temperature, the higher the average KE, the _____ the liquid evaporates

faster; (hot water is even faster than cold water)

Volatility

liquids that evaporate readily

What is the effect of IMFs on vapor pressure?

stronger the forces, the lower the vapor pressure; inverse relationship; fewer molecules will have enough KE to escape and substances with high vapor pressures are volatile, easily evaporate

What is the effect of surface area on vapor pressure?

no effect

Viscosity

resistance of a liquid to flow; molecules slide over each other

What is the effect of temperature on viscosity of a substances?

inverse relationship; viscosity decreases with increased temperature; increasing temperature increases energy and the velocity, so they interact for shorter time reducing internal friction and decreasing viscosity

What is the effect of molecular weight on viscosity of a substances?

viscosity increases with an increase in molecular weight; proportional

What is the effect of IMFs on viscosity of a substances?

the stronger the intermolecular forces, the higher the viscosity; proportional

Meniscus formation and characteristics

when adhesive forces are greater than cohesive forces, the water binds to the graduated cylinder creating a U-shaped meniscus; if the cohesive forces are greater, then the water binds to itself creating a curved downwards meniscus

Adhesive forces

bind molecules to a surface

Cohesive forces

bind molecules to each other

Surface Tension

amount of energy required to increase the surface area of a liquid

Phase Changes

Generally the heat of fusion is less than

the heat of vaporization; takes more energy to completely separate molecules, than to partially separate them

What is the term for melting?

heat of fusion

What is the term for evaporation?

the heat of vaporization

Vaporization

endothermic; liquid → gas

Melting

endothermic; solid→liquid

sublimation

endothermic; solid→gas

condensation

exothermic; gas→liquid

freezing

exothermic; liquid→solid

deposition

exothermic; gas→solid

Heating curve

plot of temperature change vs. heat added; during a phase change, adding heat causes no temperature change (equilibrium); points calculate change in H_fus and change in H_vap

Gases are liquefied by increasing __________ at some temperature

pressure

critical temperature

the minimum temperature for liquefaction of a gas using pressure; high CT means strong intermolecular forces

Critical pressure

pressure required for liquefaction

Exothermic

transfers heat to the surroundings; feels hot

Endothermic

absorbs heat from surroundings; feels cold

Coulomb’s Law

the strength of the electrostatic force (attraction/repulsion) between two charged objects; higher charges, higher electrostatic forces. longer distances, lower electrostatic forces

Metallic solids

held together by a “sea” of collectively shared electrons

Ionic soldis

sets of cations and anions mutually attracted to each other (Coulomb’s Law)

Covalent-network soldis

joined by an extensive network of covalent bonds

molecular solids

discrete molecules held together by weak forces

Metal

group of cations suspended in a sea of electrons (electron-sea model)

Alloys

materials that contain more than one element and have characteristic properties of metals; employed to change the properties of certain metals

Substitutional alloys

second element takes the place for a metal atom; homogeneus mixture; components dispersed randomly and uniformly; atoms of solid occupy positions occupied normally by a solvent atom; 2 metallic components with similar atomic radii and chemical-bonding characteristics

Interstitial alloys

second element fills a space in the lattice of metal atoms; homogeneous mixtures; components dispersed randomly and uniformly; atoms of the solute occupy positions in the “holes” between solvent atoms; solute atoms need to have smaller bonding atomic radius than solvent atoms; its element is a nonmetal that makes covalent bonds with the metal atoms; harder, stronger, less ductile

Heterogeneous alloys

components not dispersed uniformly; components are not dispersed uniformly; properties depend on the composition and manner when a solid is formed from molten mixture; formed by rapid cooling are different from slow cooling of same mixture

Properties of Ionic Solids

very high melting and boiling points; quintessential crystals; charge is centered on the anions, electronic insulators; favorable Structures with hcation-anion distances as close as possible; CsCl structure, NaCl structure and Zince blende (ZnS) structure

cesium chloride (CsCl) structure; primitive cubic lattice

two atom-basis; center atom; no lattice point inside primitive unit cell; anions sit on the lattice points at the corners and cations sit in the center of the cell; surrounded by 8 atoms;

Sodium Chloride (NaCl); rock salt structure

face-centered cubic lattice; anions sit on lattice points that lie on the corners and faces of unit cell; cations are displaced from anions along the edge of the unit cell; each cation and anion are surrounded by six ions of the opposite type; octahedral coordination environment

Zinc blende (ZnS)

face-centered cubic lattice; anions sit on the lattice points that lie on the corners and faces of the cell; cations are displaced from anions along the body diagonal of the cell; each cation and anion are surrounded by 4 of the opposite type; tetrahedral coordination geometry

When to use the ionic structure based on ion size?

cation and anion are similar in size, large coordination # is favored so CsCl structure; relative size of cation gets smaller, coordination # drops from 8 to 6, sodium chloride structure; cation size decreases a lot, coordination # reduces from 6 to 4, zinc blende structure

Weak IMFs, solubility is

low (less polarity)

Large dipole movement, solubility is

increased (greater molecular polarity)

More hydrogen bonding, solubility is

increased (greater polarity)