exam 1 semester 2

intermolecular forces (IMFs)

The various forces of attraction that may exist between the atoms and molecules of a substance. These forces hold multiple molecules of a substance together and determine many of a substance’s properties.

IMFs are a result of

electrostatic phenomena

IMFs are ____ forces

non-covalent

Intramolecular forces

the forces that hold atoms together in a molecule

the differences in the properties of a solid, liquid, or gas are a result of

differences in the strengths of the IMFs that make up each substance

the phase in which a substance exists depends on

the relative extents of its intermolecular forces and the kinetic energies of its molecules

intramolecular vs intermolecular forces

intramolecular are stronger

three types of IMFs

dispersion (london) forces, dipole-dipole attractions, hydrogen bonding

all IMFs are sometimes collectively referred to as

van der Waals forces

these substances have dispersion forces

all substances, polar or non polar

instantaneous dipoles are a result of

constant motion of electrons in molecules and atoms

an instantaneous dipole in one molecule or atom can then

distort the electrons in a neighboring atom or molecules producing an induced dipole

disperson forces

weak electrostatic forces from fleeting, temporary dipoles

stronger dispersion forces lead to

higher melting and boiling points

the strength of london forces increases with

increased number of electrons in a molecule, because larger molecules can be polarized more easily

surface area and dispersion forces

increasing surface area increases the strength of the dispersion forces

these molecules have dipole-dipole attractions

molecules with permanent dipoles (polar molecules)

polar molecules have

both dipoles dipole and dispersion, dipole dipole is stronger

hydrogen bonding is found in

certain polar molecules

hydrogen bonding

strong type of dipole-dipole, needs a H covalently bonded to O, N, or F

strong attraction of hydrogen bonding is accounted for by

the exposed nucleus for the lone pair of the adjacent molecule

water hydrogen bond

can make up to 4

ammonia hydrogen bond

can make up to 2

why is ice less dense than water

water molecules in their crystal form have an open-structured, six-sided arrangement. as a result, water expands upon freezing

liquids exhibit

low compressibility, lack of rigidity, and high density

three properties unique to liquids

viscosity, surface tension, capillary action

viscosity

measure of a liquid’s resistance to flow, measured by the rate in which a metal ball passes through the liquid

viscosity and IMFs

liquids with large IMFs or molecular complexity tend to be highly viscous

temperature and viscosity

viscosity of the liquid decreases

adhesive forces

force of a attraction between molecules of different chemical identities

cohesive forces

force of attraction between identical molecules

surface tension

the energy required to increase the surface of a liquid

surface tension results from

cohesive forces between molecules at the surface of a liquid

surface tension and IMFs

large IMFs have high surface tension

capillary action

flow of liquid within a porous material due to the attraction of the liquid molecules to the surface of the material and to other liquid molecules

capillary action is exhibited by

polar liquids and involves spontaneous rising of a liquid in a narrow tube

concave mensiscus

this occurs when the adhesive forces of the liquid and glass are greater than the cohesive forces of the liquid

convex mensicus

this occurs when the adhesive forces of the liquid and glass are less than the cohesive forces of the liquid

capillary action is due to

adhesive forces, based o strength of the adhesive forces, the liquid may rise or fall

extent of the rise or fall and surface tension

directly proportional

extent of the rise and fall and density of the liquid/radius of the tube

inversely proportional

thin layer chromatography

uses capillary action in which a layer of liquid is used to separate mixtures from substances

liquids are constantly

vaporizing, at a given instant some of the liquids have enough energy to escape the liquid phase

in an open container all liquid will

eventually evaporate

if a liquid is in a closed vessel with space above it

a partial pressure of the vapor state builds up in this space

when the rate at which the liquid vaporizes is equal to the rate at which the vapor condenses

a dynamic equilibrium is established

after equilibrium has been reached

the number of molecules in the gas phase does not change with time

vapor pressure

the pressure of a vapor in equilibrium with a liquid

the vapor pressure of a substance is independent

of the volume of the container

vapor pressure and temperature

vapor pressure increases with increasing temperature

boiling point

when the vapor pressure of a liquid increases enough to equal the external atmospheric pressure

boiling

as the temperature of a liquid increases, the vapor pressure increases until it reaches atmospheric pressure, temperature remains constant

normal boiling point

1 atm, 101.3 kPa, 760 torr

the lower the vapor pressure of a substance

the higher the boiling point

boiling point and atmospheric pressure

at lower pressures, the boiling point lowers

vaporization

an endothermic process

enthalpy of vaporization

the energy change associated with the vaporization process

condensation

is an exothermic process

enthalpy of condensation

the energy associated with the condensation process

enthalpy of fusion

endothermic, melting

enthalpy of freezing

exothermic, freezing

temperature remains constant

during phase change

sublimation

solid to gas state, bypassing the liquid state

deposition

gas into solid, bypassing the liquid state

sum of enthalpy of fusion and vaporization equals

enthalpy of sublimation

heating curve

a plot of temperature against time for a process where energy is added at a constant rate

heating curves depict

changes in temperature that result as the substance absorbs increasing amounts of heat

plateaus in the curve

are exhibited when the substance undergoes phase transitions

phases of a heating curve

ice, ice and water, water, water and steam, steam

heat required to increase the temperature of a substance

q = mass times specific heat times temperature change

heat absorbed or released during phase change

q = moles x change in H

triple point

the temperature at which all three phases exist simultaneously

critical point

the critical pressure and critical temperature, together, define this point

critical pressure

pressure required to produce liquefaction at critical temperature

critical temperature

the temperature above which a liquid cannot be liquefied, irrespective of pressure applied

supercritical fluid

a state which has a density characteristic of a liquid but the flow properties of a gas allowing it to diffuse through substances easily

the slope/liquid boundary is a negative slope

melting point of ice decreases with external pressure

regelation

refreezing of water derived from the melting of ice under pressure when the pressure is relieved

crystalline solids

regular arrangement, position of the components represented by a latice, the smallest repeating unit is called a unit cell

amorphous solids

disorder in structure, may undergo a transition to the crystalline state under appropriate conditions, no unit cell

types of crystalline solids

atomic solids, ionic solids, molecular solids

atomic solids

have atoms at the lattice points that describe its structure, metallic, network, and group 8A solid

ionic solids

possess ions at the points of the lattice that describe their structures

molecular solids

have discrete covalently bonded molecules at each lattice point

what solids are these

metallic solids, formed by metal atoms, uniform distribution of atomic nuclei within a sea of delocalized electrons, strong and nondirectional (easy to move metal atoms, hard to separate them) malleable and hard to shatter

alloy

a substance that contains a mixture of elements and has metallic properties

what is this

substitutional alloy, some of the host metal atoms are replaced by other metal atoms of similar size

what is this

interstitial alloy, some of the holes in the closest packed metal structure are occupied by small atoms

networks solids

atomic solids that are held together by directional covalent bonds, which form solids that can be viewed as giant molecules

properties of network solids

brittle nature, hard, high melting point, ineffective conductors of heat and electricity

diamond properties

network solids, pure carbon, six membered rings, tetrahedral bonding, three-dimensional network, denser than graphite, poor electrical conductor, hard substance, good heat conductor, transparent

graphite

pure carbon, planar network solid, six membered carbon ring, each carbon atom is bonded to three others, conducts heat and electricity, strong interaction within the planes, weak interaction between planes, slippery and black

graphene

pure carbon, thinnest compound at 1 atom, exception strength due to carbon-carbon valence bond

graphene sheets can be formed into

buckballs (fullerenes), nanotubes, and stacked layers

carbon allotropes are

strong, lightweight, good conductors of electricity and heat

AM III

yellow tint, made of carbon, very hard, semi conductor, has ordered structure when viewed closely but highly disorded with less magnificatio

group 18 solids

gases, full outer shell and usually unreactive, when cooled or compressed they can become solid

ionic solids

consist of ions at the lattice points held together by electrostatic interactions, high melting points, hard, brittle, and shatter, do not conduct electricity unless they are molten or dissolved

molecular solids

neutral molecules with strong covalent bonding but weak IMFs

how to determine packing efficiency

count the number of atoms in the cell, coordination number is the number of other atoms touching the atom under consideration