AP Chemistry Unit 3

intermolecular forces

intermolecular forces

forces of attraction between molecules (not chemical bonds)

van der Waals forces

intermolecular forces of attraction

dipole-dipole interactions

• molecules have permanent dipoles attracted to one another

• positive end of one attached to negative end of the other

• these forces matter when molecules are close to each other

• more polar -> higher boiling point

permanent dipole

permanent separation of electrical charge in a molecule due to unequal distributions of bonding and/or lone pairs of electrons

hydrogen bonding

• H is bonded to N,O, or F

• bonding caused by high electronegativity and small size of N, O, and F

London dispersion forces

• attractions between an instantaneous dipole and an induced dipole

• present in all covalent molecules, weakest bond

instantaneous dipole

temporary dipole that occurs for a brief moment in time when the electrons of an atom or molecule are distributed asymmetrically

induced dipole

a dipole temporarily created in an otherwise nonpolar molecule, induced by a neighboring charge

polarizability

the ease with which the electron distribution in the atom or molecule can be distorted

viscosity

a liquid's resistance to flow (related to ease with which molecules move pass each other)

surface tension

the force that acts on the surface of a liquid and that tends to minimize the area of the surface (caused by attraction between liquid molecules)

melting point

the temperature at which a substance changes from a solid to a liquid

boiling point

the temperature at which a substance changes from a liquid to a gas

cohesion

an attraction between molecules of the same substance

adhesion

an attraction between molecules of different substances

capillary action

• the spontaneous rising of a liquid in a tube

• happens when a liquid is strongly attracted to the tube

• ex: water climbs through plant's xylem

vapor pressure

the pressure caused by the collisions of particles in a vapor with the walls of a container

amorphous solids

the particles are not arranged in a regular pattern

crystalline solids

highly regular arrangement of their components (can be any repeating structure, not just cubic)

lattice

a three-dimensional system of points designating the positions of the centers of the components of a solid (atoms, ions, or molecules)

ionic solids

• ion-ion interactions are the strongest

• lattice points occupied by ions

• held together by electrostatic attraction

• hard, brittle, high melting point

• poor conductor of heat and electricity

network atomic solids

• stronger than IMFs but weaker than ion-ion

• lattice points occupied by atoms

• held together by covalent bonds

• hard, high melting point

• poor conductor of heat and electricity

metallic atomic solids

• weaker than covalent bonds, but can be in the low end of covalent bonding

• lattice points occupied by metal atoms

• held together with metallic bonds

• soft to hard, low to high melting points

• good conductors of heat and electricity

molecular crystals

• lattice points occupied by molecules

• held together by IMFs

• soft, low melting point (lowest)

• poor conductor of heat and electricity

network solids

• diamond/graphite

• silicon dioxide/nitride

• boron nitride

pressure equation

P=force/area

barometer

measures air pressure

Boyle's Law

P1V1=P2V2

Charles' Law

V1/T1=V2/T2

Gay-Lussac's Law

P1/T1=P2/T2

Combined Gas Law

P1V1/T1=P2V2/T2

Avogadro's Law

V1/n1=V2/n2

Ideal Gas Law

PV=nRT

Dalton's Law of Partial Pressures

the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture

mole fraction

ratio of the number of moles of a given component in a mixture to the total number of moles in the mixture (does not change with temp)

kinetic molecular theory of gases

a model used to explain the behavior of (ideal) gases

• volume of individual particles is negligible because particles are so small

• particles are in constant motion (cause pressure)

• gas molecules do not exert attractive or repulsive forces on each other

Maxwell-Boltzmann distribution

shows the spread of energies that molecules of gas or liquid have at a particular temperature

diffusion

the process by which molecules move from an area of higher concentration to an area of lower concentration

effusion

a process by which gas particles pass through a tiny opening into a chamber

Graham's law of effusion

• states that the rate of effusion for a gas is inversely proportional to the square root of its molar mass

Henry's Law

the solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on the surface of the liquid

solution

a homogeneous mixture of two or more substances

solute

the substance that is dissolved; present in smaller amounts

solvent

the substance in which the solute dissolves; present in larger amounts

miscible

liquids that dissolve freely in one another in any proportion

ion-dipole

• ionic compounds dissolve in polar solvents

dispersion forces

• non-polar solids dissolve in non-polar solvents

hydration of ionic compounds

• positive ends of H2O attract to anions; negative ends of H2O attract to cations

• hydation of ions causes salt to fall apart in water

electrolyte

a substance that dissolves in water to give a solution that conducts electric current

nonelectrolyte

a substance that dissolves in water to give a solution that does not conduct an electric current

saturated solution

contains the maximum amount of solute for a given quantity of solvent at a constant temperature and pressure

unsaturated solution

a mixture that contains less dissolved solute than is possible at a given temperature

supersaturated solution

contains more dissolved solute than a saturated solution at the same temperature

molarity

the number of moles of solute per liter of solution (changes with temp)

dilution

• the process of adding solvent to lower the concentration of solute in a solution

• M1V1=M2V2

three interactions in solution process

• solute-solute interaction (separate solute)

• solvent-solvent interaction (overcome IMFs)

• solvent-solute interaction

chromatography

separates chemical species using differing strengths of IMFs between and among solution and surface of stationary phase

mobile phase of chromatography

solvent

stationary phase of chromatography

solid, chromatography paper

fractional distillation

separation of liquid mixture by vaporization and then condensing vapor into liquid

light equation

c = λv

Planck's equation

E=hv

photoelectric effect

the emission of electrons from a metal when light shines on the metal

electromagnetic spectrum

the range of wavelengths or frequencies over which electromagnetic radiation extends.

spectroscopy

the study of the interaction of electromagnetic radiation with matter

absorption spectroscopy

measures the amount of light absorbed by the sample as a function of wavelength

UV and visible light cause

• electronic transitions within atoms

• can be used to gather information about electronic configurations

infrared radiations cause covalent bonds to

• bend, stretch, and vibrate (depending on bond type and functional group)

• can be used to distinguish between compounds having different types of bonds

microwaves cause

• molecular rotations

• can determine location of different atoms in a molecule

• give information about the chemical composition and structure of molecules

spectrophotometer

an instrument that measures the absorbance or transmittance of light, as a function of wavelength

colorimeter

a device that measures the color of foods in terms of value, hue, and chroma

cuvette

a straight-sided, optically clear container for holding liquid samples in a spectrophotometer or other instrument. (1 cm thick)

absorbance

the amount of light absorbed by the sample

transmittance

the amount of light that passes through the sample

Beer's Law

• explains the relationship between absorbance at a given wavelength and concentration

• A = εbc