Gen. Chem Ch. 8-9 Exam

Change of State

the change of a substance from one state of matter (gas, liquid, or solid) to another

Three phases of matter

solid, liquid, and gas; depends on the relative strength of the attractive forces between particles compared with the kinetic energy of the particles

Kinetic Energy

energy associated with motion and is related to the temperature of a substance

Temperature

Measure of the kinetic energy associated with a substance or object

Temperature vs. kinetic energies

As temperature increases, kinetic energies of the particles also increases

Intermolecular forces

These associated with particles are independent of temperature; a constant quantity

Gases vs. kinetic energy

This energy of the particles has become much greater than teh attractive forces between the particles

Gas particles

move about freely, are far apart from one another, have almost no influence on a neighboring particle

Liquids vs. kinetic energies

these energies are only large enough to allow the particles to move around relative to each but not to escape (get away) from each other; this energy of the particles are roughly the same as the forces of attraction (intermolecular forces) between particles

Solids vs. kinetic energies

Attractive forces between particles are much greater than this type of energy of the particles; the particles are held together in a rigid pattern and can only vibrate

Characteristics of change of state

Every change of state is reversible and can be explained by changes in enthalpy and entropy

Enthalpy change

deltaH, measure of the heat absorbed or released during a given change of state

Entropy change

deltaS, a measure of the change in molecular disorder or freedom that occurs during a process

Melting of a solid to liquid

Heat is absorbed and enthalpy change is positive (endothermic); as heat is absorbed, the kinetic energy of molecules increases until it is sufficient to overcome the forces of attraction

Freezing of a liquid to a solid

Heat is released and enthalpy change is negative (exothermic); the potential energy of attractive forces between molecules is converted to thermal energy

Amount of heat required to melt a mole of ice vs. freeze a mole of wayer

equal for both processes and is equal to 6.02 kJ

Entropy vs. Melting

Disorder increases because particles gain freedom of motion and change in entropy is positive

Freezing vs. Entropy

disorder decreases as particles are locked in position and entropy change is negative

General rules for entropy

for any substance, its solid form is more ordered than the liquid state and liquid state has lower entropy than gaseous version; solid form has lower entropy than liquid

General rules for reactions

processes lower in energy are favored to happen, exothermic processes result in a process ending in lower energy state, higher entropy results in lower energy; negative enthalpy change an positive entropy change favors a process being spontaneous

Enthalpy and Entropy vs. Phase changes

Contrary; melting is unfavored by positive enthalpy change but favored by positive entropy change, freezing is favored by negative enthalpy change but unfavored by negative entropy change

Enthalpy and Entropy balance

Exact temperature in which these two factors balance out is the temperature where the phases are in equilibrium with one another

Melting point

temperature at which the liquid phase is in equilibrium with the solid phase

Boiling point

temperature at which the gas phase is in equilibrium with the liquid phase

Intermolecular forces

forces that act between molecules or discrete atoms and hold them close to one another; also called van der Waals forces

Intermolecular forces in gases

These forces are negligible in this state

Intermolecular forces in liquids and solids

the stronger the forces, the higher the melting and boiling points are

Types of intermolecular forces

dipole-dipole, London dispersion, and hydrogen bonding

Dispersion forces

temporary polarity in the molecules due to unequal electron distribution leads to attractions; also known as London forces, induces a dipole in surrounding molecules, all molecules and atoms have this

Dipole-dipole attractions

permanent polarity in the molecules due to their structure leads to attractive forces

Hydrogen bonds (H-bonding)

an especially strong dipole-dipole attraction (stronger than other types) results when H is attached to an extremely electronegative atom; substances that have these have higher BP and MP, not nearly as strong as chemical bonds (2-5% strength of covalent bonds)

London dispersion forces

All molecules experience these forces, at any given instant there may be more electrons at one end of the molecule than at the other creating a temporary dipole, induces a dipole in neighboring particle; the larger the molecular weight and surface area the greater the temporary polarization of a molecule

Temporary dipole

results from fluctuations in the electron distribution in atoms and molecules; region with excess electron density has partial negative charge, region with depleted electron density has partial positive charge

Noble gases

ALL are nonpolar atomic elements

Boiling point vs. attractive forces

Stronger the forces, the higher the point will be

Molar mass vs. strength of forces

As this value increases, the number of electrons increases and the strength increases; easier to distort the electron cloud making bigger temporary dipoles

Molecular shape vs. intermolecular force strength

more oblong shaped molecules have more places for interaction and electron clouds more easily distorted, stronger forces; more spherically shaped molecules are more difficult to distort the electron clouds, weaker forces

London dispersion forces rules

The larger the molecular weight and SA, greater the temporary polarization; branched chains (spherical) have lower boiling points than straight chains since straight chains have more surface-to-surface contact

Dipole-dipole forces rules

molecules with polar covalent bonds may have a net molecular polarity, positive and negative ends of different molecules are attracted to one another; polar molecules have permanent dipole, adds attractive forces between molecules which raises BP and MP relative to nonpolar molecules with similar size and shape

Hydrogen bonding rules

very electronegative atom is bonded to hydrogen and strongly pulls bonding electrons to it; because hydrogen has no other electrons its electrons pull away and nucleus becomes desheilded, exposed proton is very strong center of positive charge and attract all electron clouds from neighboring molecules

Hydrogen bonding with O, N, or F

unusually strong dipole-dipole interaction; in water, there are four of these bonds due to two lone pairs on oxygen atoms and two hydrogen atoms

Kinetic-molecular theory of gases

Group of assumptions that explain the behavior of gases

Postulates of Kinetic-molecular theory of gases

1) a gas consists of particles moving randomly with no attractive forces; different gases mix together in all proportions and relatively quickly as a result

2) amount of space occupied by gas particles themselves is much smaller than amount of space between particles; volume taken up is empty space which accounts for compressibility and low densities

3) Average kinetic energy of gas particles is proportional to Kelvin temperature; gas have more and move faster as temperature increases

4) collisions of gas particles either with other or with wall of container are elastic; total kinetic energy is constant

Ideal gas

gas that obeys all assumptions of the kinetic-molecular theory

Pressure (P)

force per unit area pushing against a surface; result of gas hitting walls of vessel

Atmospheric pressure

column of air weighing 14.7 lbs presses down on each square inch of Earth’s surface at sea level

Mercury barometer

measures atmospheric pressure by determining the height of mercury column in sealed glass tube

mmHg or torr

common unit of pressure after the height of a column of mercury in a barometer

Manometer

used to measure gas pressure; difference between mercury levels indicates the difference between gas and atmospheric pressures

SI Unit for Pressure

Pascal (Pa)

Standard pressure unit

In chemistry it is the atmosphere (atm)

Conversions of atm

1 Atm = 760 mmHg, 1 Atm = 14.7 psi, 1 Atm = 101, 325 Pa

Gas laws

Series of laws that predict the influence of pressure (P), volume (V), and temperature (T) on any gas or gas mixture

Simplest state of matter

Gases; only need to know mass, volume, pressure and temperature to determine qualities about the gas and molar mass and don’t need to know the identity to get these values

Robert Boyle

First person to study the properties of gases in the 1600s, studied how a sample of gas (constant mass and temperature) was changed by differing the volume of the gas occupied

Boyle’s findings (volume and pressure)

Volume of a gas decreases proportionately as its pressure increases at constant n and T (inversly related); if the pressure of a gas is doubled, the volume is halved

Boyle’s Law

The volume of a gas is inversely proportional to its pressure for a fixed amount of gas at a constant temperature; P times V is constant when the amount of gas (n) and temperature (T) are kept constant, because PV is constant value for fixed n and T the starting PV must equal the ending PV

Kinetic molecular theory and Boyle’s Law

if sample size and temperature are kept constant, a decrease in volume should result in the particles hitting the walls of the container more often thereby increasing the pressure of the gas in the vessel

Charles’s findings

The volume of a gas is directly proportional to its kelvin temperature at constant n and P (if the kelvin temp is doubled, the volume is doubled); volume is directly proportional to kelvin temperature

Charles’s Law

Volume of a gas is directly proportional to its kelvin temperature for a fixed amount of gas at constant pressure; V divided by T is constant when n and P are held constant (V1/T1 = V2/T2)

Kinetic molecular theory and Charles’s Law

if pressure and sample size are constant, heating a gas increases the kinetic energy of the particles (1/2 mv2); this will result in the particles hitting the walls of the containing more often and so to keep the pressure constant the volume must expand

Gay-Lussac’s Findings

The pressure of a gas is directly proportional to its kelvin temperature at constant n and V; if the kelvin temperature is doubled, the pressure doubles

Gay-Lussac’s Law

The pressure of a gas is directly proportional to its kelvin temperature for a fixed amount of gas at a constant volume; P divided by T is constant when n and V are held constant (P1/T1 = P2/T2)

Combined gas law

For a fixed amount of gas, Boyle’s law, Charles’s Law, and Gay-Lussac’s law can be merged (P1V1/T1 = P2V2/T2)

Avogadro’s Law

The volume of a gas is directly proportional to its molar amount at a constant pressure and temperature; V is divided by n is constant when P and T are held constant (V2/V1 = n2/n1)

Explanation for Avogadro’s Law

because particles in gas are so tiny compared to empty space there is no interaction as proposed by kinetic molecular theory; chemical identity of particles does not matter and the value of the constant k in the equation V/n = k is the same for ALL gases

Avogadro’s Law and T and P

Values of temperature and pressure do not matter; only necessary for them to be the same for both gases but convenient to define set of conditions called STP

STP

Standard temperature is 273 K (0ºC) and standard pressure is 1 atm

Standard molar volume

Of any ideal gas at STP is 22.4 L/mol

Ideal gas law

relationships among four variables P, V, T, and n for gases can be combined; PV/nT = R OR PV = nRT

Gas Constant

Constant R; its value depends on units chosen for pressure, R = 0.0821 L(atm)/mol(K) OR R = 62.4 L(mmHg)/mol(K)

Derived laws of gases

PV = nRT can be rewritten as PV/nRT = 1; if we have multiple set of conditions we could get P1V1/n1RT1 = 1 and P2V2/n2RT2 = 1

Rearranging derivative of gas laws

P2V2/P1V1 = n2RT2/n1RT1

Derivative of Boyle’s Law

P2V2/P1V1 = 1 or P2V2 = P1V1

For any sample of any gas at constant temp

The product of PV should be a constant value and should be a constant graph

Partial Pressure

the contribution of a given gas in a mixture to the total pressure; mixtures of gases behave the same as pure gases and obey the same laws, pressure exerted by each gas depends on frequency of collisions of its molecules with walls of the container (does not change with other gases)

Dalton’s Law

the total pressure exerted by a gas is the sum of the partial pressures of the components in the mixture; P(total) = P(gas 1) + P(gas 2) + P(gas 3)

Determining pressure

Only the total number of particles needs to be known to determine the pressure since the identity does not matter (ex: 10 particles of 5 different gases will exert the same pressure as 50 particles of a single gas)

Pressure exerted by single gas in mixture

Directly proportional to its fraction of fraction of particles in the mixture (ex: if 10% of the particles in a mixture are gas A, gas A exerts 10% of the total pressure)

Vapor

The gas molecules are in equilibrium with a liquid; molecules are in constant motion in liquid state and if a molecule is near the surface and has enough energy it can break free and escape into the gas state

Open containers and vapors

Gaseous molecules move away from container and process will continue until all the liquid molecules evaporated (ex: puddles evaporate despite not being near BP)

Closed containers and vapors

random motion of molecules occasionally brings them back into the liquid; at beginning of this process, it is less likely for molecules to return to liquid state but as more evaporates there are more chances for gas molecules to return to liquid phase

Vapor pressure

partial pressure of vapor molecules in equilibrium with a liquid; when evaporation and condensation take place at same rate and concentration of vapor becomes constant

Factors influencing vapor pressure

Temperature, identity of liquid, and intermolecular forces; strong IM forces causes higher MP and BP and the stronger IMFs the lower the vapor pressure at a given temperature

Temperature vs. Vapor pressure

Increasing temperature increases VP; as temp rises the molecules have greater kinetic energy and this causes more molecules to escape to the gas phase, VP continues to rise until all molecules have kinetic energy to vaporize and occurs when VP and atmospheric pressure are equal

Normal boiling point

boiling point at a pressure of exactly 1 atmosphere

Factors affected BP

intermolecular forces present in each liquid and altitude

Viscosity

measure of a liquid’s resistance to flow; increases with increasing IMFs

Surface tension

caused by the difference between IMFs experienced by molecules at the surface of the liquid and those experienced by molecules in the interior

Crystalline solid

one whose atoms, molecules, or ions are rigidly held in an ordered arrangement; can be further categorized as ionic solids, molecular solids, covalent network solids

Ionic solids

crystalline solids hose constituent particles are ions; a crystal composed of alternating positive and negative ions in a regular three-dimensional arrangement held together by ionic bonds

molecular solids

crystalline solids whose constituent particles are held together by IMFs (ex: ice)

Covalent network solids

crystalline solids where atoms are linked together by covalent bonds in a giant, three-dimensional array; in effect, it is one very large molecule (ex: diamond)

Metallic solids

can be viewed as vast three-dimensional arrays of metal cations immersed in sea of electrons; electron sea acts as glue to hold cations and mobile carrier of charge to conduct electricity, bonding attractions extend uniformly in all directions (makes metals malleable rather than brittle), when metal crystal receives sharp blow the electron sea adjusts to new distribution of cations

Amorphous solid

one whose constituent particles are randomly arranged and have no ordered long-range structure; often result when liquids cool before they can achieve internal order or molecules are large and tangled together, soften over wide temperature range and shatter to give pieces with curved rather than planar faces (ex: glass, tar, opal, hard candies)

Heat of fusion

quantity of heat required to completely melt one gram of a substance once it has reached its melting point

Heat of vaporization

quantity of heat needed to completely vaporize a liquid at its boiling point

IMFs vs. heats of fusion and vaporization

stronger they are leds to greater heats of these

Energy needed for phase change

Depends only on the amount of the substance and the heat of fusion and vaporization

How to determine how much is required to geat or cool a phase

Heat = (mass x specific heat x temp change)