CHEM 1212 UGA Exam 1

Fusion (melting)

solid to liquid

energy is absorbed.

vaporization

liquid to gas

energy is absorbed

sublimation

solid to gas

energy is absorbed

freezing

liquid to solid

energy is released

condensation

gas to liquid

energy is released

Deposition

gas to solid

energy is released

Relationship between IMFs and Enthalpy of vaporization

As the strength of the IMFs in a series of liquids increases, the enthalpy of vaporization values for the liquids increase.

Relationship between IMFs and Boiling point and Surface tension

Both normal boiling point and surface tension at a given temperature tend to increase as intermolecular forces increase.

Relationship between viscosity and IMFs

The one with the higher viscosity is the one with the stronger intermolecular forces and therefore, would be expected to be the one with the higher normal boiling point and the higher surface tension.

electrostatic forces

Forces of attraction or repulsion caused by electric charges

intermolecular forces

Interactions between molecules, between ions, or between molecules and ions

gas

IMFs- Generally weak

Compressibility- High

Takes on shape and volume of container

Ability to flow- High

Liquid

IMFs- generally intermediate

Compressibility- very low

Takes on shape of container; volume limited by surface area

Ability to flow- moderate

solid

IMFs- generally strong

compressibility- almost none

Maintains own shape and volume

ability to flow: almost

The stronger the IMFs between particles

the harder it is to separate the particles and the more likely the substance is a solid

When the IMFs between particles are weak

the substance made up of those particles is most likely a gas

Enthalpy of vaporization

The quantity of energy as heat required to convert one mole of a liquid to a gas at constant temperature

relative strength of IMFs in two hydrocarbons

the greater the enthalpy of vaporization the stronger the IMFs

vapor phase

the vapor pressure in the flask increases as molecules at the surface of the liquid escape into the gas phase

IMFs as temperature increases

-more molecules have the minimum kinetic energy needed to enter the gas phase,

-the number of molecules in the gas phase increases, and

-vapor pressure increases.

IMFs and vapor pressure for liquids

As the strength of the IMFs in a series of liquids increases, the vapor pressures of the liquids decrease.

For any liquid, as the temperature increases,

the vapor pressure increases because of the increasing fraction of molecules above the escape threshold.

IMFs and boiling points for liquids

As the strength of the IMFs in a series of liquids increases, the boiling points of the liquids increase.

Clausius-Clapeyron equation

The mathematical relationship between vapor pressure (P), temperature (T), and the strength of IMFs

second form of the Clausius-Clapeyron

(DOUBLE LN)

used to calculate the vapor pressure of a liquid at a given temperature if both the enthalpy of vaporization of the liquid and the vapor pressure at another temperature are known.

bulkscale forces

ohesive forces are the result of IMFs, the molecular-level interactions in a liquid

viscosity

the resistance of a liquid to flow

relationship between IMFs and viscosity

strong IMFs result in high viscosity for a liquid

Dipole-dipole intermolecular forces

The electrostatic force between two neutral molecules that have permanent dipole moments. They occur between two polar molecules

Dipole-dipole IMFs

are generally weaker than the ionic and covalent forces in ionic solids and covalently bonded compounds, respectively.

-occurs between polar molecules

Hydrogen bonding

Attraction between a hydrogen atom and a very electronegative atom to produce an unusually strong dipole-dipole attraction

-an unusually strong type of dipole-dipole IMF that occurs between molecules with H-N, H-O, or H-F bonds.

Dipole-induced dipole intermolecular forces

The electrostatic force between two neutral molecules, one having a permanent dipole and the other having an induced dipole

-occur between polar and non polar molecules

-are generally weaker than dipole-dipole forces, the magnitude of the induced dipole can result in these forces being stronger than dipole-dipole forces

induced dipole

Separation of charge in a normally nonpolar molecule, caused by the approach of a polar molecule

-a created temporary dipoles for non polar molecules

polarizability

Polarizability increases as the number of electrons in a molecule increases, and therefore, it increases with increasing molar mass and molecular size.

Induced dipole-induced dipole intermolecular forces

are attractive forces that occur between nonpolar molecules.

London dispersion forces

*are present in all molecules*

The attractive forces between these temporary dipoles, induced dipole-induced dipole IMFs

IMFs attractions between two neutral molecules, both having induced dipoles

-generally weak, but can be stronger than dipole-dipole forces when they occur between highly polarizable molecules

soluble

A compound that dissolves in a solvent to an appreciable extent

insoluble

A term describing a compound that does not dissolve to an appreciable extent in a specific solvent

saturated

A solution in which the solute concentration is at the solubility limit at a given temperature

- in order words, only a certain amount of g will dissolve. if added more some will not dissolve

unsaturated

A solution in which the solute concentration is less than the solubility limit at a given temperature

- if you add an amount, and I completed dissolves, and then you add more and that completely dissolves too, then the first amount is unsaturated bc it could use /need more

supersaturated

A solution in which the solute concentration is greater than the solubility limit at a given temperature

- a supersaturated sodium acetate solution can be prepared by dissolving a large quantity of the solid in a small amount of warm water and then allowing the solution to cool slowly to room temperature. A slight disturbance such as the introduction of a small particle of solid or simply bumping the solution container will result in precipitation of solid sodium acetate from the solution and a saturated sodium acetate solution. (having too much)

miscible

A term used to describe two liquids that intermix completely

-they form a homogeneous solution

immiscible

A term used to describe two liquids that do not intermix

-form heterogenous mixture

Molarity

moles of solute/liters of solution

weight percent

mass of solute/mass of solution x 100

molality

The number of moles of solute per kilogram of solvent

mole fraction (x)

the composition of any solution, however, not only one that is a mixture of gases.

amount (mol) component A / total amount (mol) in solution

parts per million

When a solution contains a very low amount of solute, its concentration is often reported in units of parts per million

- The mass of solute (g) in 10^6 g of solution

-mg/kg

parts per billion

the mass of solute (g) in 10^9 g of solution

-micro-gram/kg

crystalline solid

a solid in which the particles are arranged in a regular way

ex- diamond , table salt NaCl and Sugar

amorphous solid

A solid that lacks long-range regular structure and displays a melting range instead of a specific melting point

-the particles that make up the solid are arranged in an irregular manner

ex:synthetic fibers, plastics, and glasses, but pure solid substances, such as elemental phosphorus or sulfur

ionic solids

a solid formed by the condensation of anions and cations

covalent solids

A crystalline solid consisting of a three dimensional extended network of atoms held together by covalent bonds

cubic unit cell

a=b=c

=90 degrees

molecular

-dipole-dipole

-hydrogen bonds

-LDF

ex- CO2

diagram = 3 atom bonds bonded together with few spaces

melting point- moderate to low

ionic (ions)

ion-ion

ex- NaCl

diagram= big and small molecules bonded with space between them

melting point- high to very high

covalent (covalent network)

-covalent bonds

-diamond

diagram= small atoms connected through lines

melting point- very high

metallic

-metallic bonds

-Na

diagram= small atoms tightly packed

melting point- variable

radius in pm

a= 4r/sqrt3

face-centered cubic (FCC) unit cell

A cubic unit cell in which there are lattice points on each corner of the cube and a lattice point centered on each cube face

entropy

is a measure of the dissipated energy within a system that is unavailable to do work at a given temperature.

temperature and entropy

as temperature increases, molecular motion increases and so does entropy.

entropy is linked to

the motion (or mobility) of atoms and molecules

the more freedom of motion the particles have

the higher the entropy

endothermic

if the IMFs in the pure substance are stronger than those formed between species in the solution, the overall enthalpy change for the mixing process is positive

exothermic

if the IMFs in the mixture are stronger than those in separated species, the enthalpy change for the mixing process is negative

a process is favored by entropy and enthalpy if

-any chemical process that increases the free motion of molecules

-those that result in the formation of stronger IMFs

Positive enthalpy change (reduction in strength of IMFs) and Positive energy change (increase in mobility of particles)

favored in higher temperatures

Negative enthalpy change (increase in strength of IMFs) and Positive energy change (increase in mobility of particles)

Favored at all temperatures

Positive enthalpy change (reduction in strength of IMFs) and Negative entropy change (Reduction in mobility of particles)

disfavored at all tempertaures

Negative enthalpy change (increase in strength of IMFs) and Negative entropy change (Reduction in mobility of particles)

favored at lower temperatures

evaportation involves breaking bonds that hold one molecule near another. Bond breaking requires energy

so this process is called endothermic

Bond breaking is

disfavored in terms of enthalpy change

Separating molecules in the liquid to form the vapor is disfavored in terms of enthalpy change. So why does evaporation occur? The answer lies in the entropy change that also occurs. In which state do the molecules of the compound experience greater freedom of movement?

The vapor state is favored in terms of entropy because the molecules have complete freedom of movement. Greater freedom of movement means greater entropy, which is favored for a system

Gases have

no appreciable IMFs

molecules are broken

-endothermic

-enthalpy disfavored

molecules are formed

-exothermic

-enthalpy favored

All hydrocarbons are

non polar molecules

polar liquids are miscible in

polar liquids

non polar liquids are miscible in

non polar liquids

water and hydrocarbon

immicible

ion-dipole intermolecular force

attractive forces that occur between ions and polar molecules

when ions separate from one another (bond breaking)

requires energy input so its endothermic and disfavored

enthalpy of hydration

is the enthalpy change when one mole of a gaseous ion dissolves in water, forming a hydrated ion

the enthalpy of hydration is an

exothermic process and is thus generally favored

As ion size increases (and ion charge is held constant)

the two species attracted to each other are farther apart and the enthalpy of hydration decreases

As ion charge increases (and ion size is held constant)

the attractive force increases and so does the enthalpy of hydration

ions with small radii and large charge magnitude have the greatest (most negative)

enthalpies of hydration

Hydration number increases with increasing charge

and decreasing ionic radius

Ion & Radius

Na+ - 116pm

Rb+ - 166pm

Cd 2+ - 109pm

-largest magnitude enthalpy of hydration = Cd 2+

-largest hydration number = Cd 2+

-entropy-favored to dissolve = Rb+

Of the ions given, Cd2+ has both the highest charge and the smallest radius. Therefore, it has the largest magnitude hydration enthalpy. Hydration number also follows the same trends, and Cd2+ will have the largest hydration number.

During the dissolution process, water molecules from the solvent are constrained to be in the hydration sphere of the cation. Therefore, in terms of entropy of dissolution, the smaller the hydration number, the more favored the dissolution. Of the three, Rb+ is the ion with the lowest charge and the largest radius, so it will have the smallest hydration number and will be the most favored to dissolve in terms of entropy.

Hydration enthalpy increases in magnitude as the

ion charge increases and as the radius increases

Hydration number increases as the

ion charge increase and the radius decreases

The smaller the hydration number, the

more favored the dissolution in terms of entropy

If the pressure in the flask is increased, either by introducing more gas into the container or decreasing the volume of the container,

the solubility of the gas in the liquid increases

As pressure is increased, there are more collisions between the gas molecules and the solvent; thus the rate of dissolution

increases

Henry's law

the solubility of a gas is proportional to the pressure of the gas above the solution

S= kH x P

S= solubility

kH= Henry's law constant for a specific solute, solvent, and temperature

P= partial pressure of the solute gas

As temperature increases, molecular motion

increases and so does entropy.

An increase in temperature always makes the change in entropy larger, making the solution process more favorable.

Gas dissolution processes are generally exothermic because

new IMFs are formed between the gas and solvent molecules but no IMFs are broken between the gas molecules.

IMFs between solvent molecules are broken

-endothermic

-enthalpy disfavored

New IMFs between gas and solvent molecules are formed.

-exothermic

-enthalpy favored