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When a system goes through a change in one or more of its properties

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Process

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Delta U = Q - W where Delta U is change in system's internal energy, Q is heat added to system and W is work done by the system

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First law of thermodynamics

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Occur at constant temp, therefore U is constant. Therefore Delta U is 0 and Q = W

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Isothermal processes

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Total internal energy of system directly related to temperature

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Why is isothermal imply no Delta U?

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Hyperbola (looks like exponential decay) where area under curve = work AND ALSO the heat that entered the system

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Isothermal process on PV graph

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No heat exchange. Therefore Delta U = -W Therefore Delta U equals work done ON the system Temperature not constant.

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Adiabatic process

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Pressure is constant. Flat line on P-V graph

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Isobaric process

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aka isochoric; no change in volume therefore no work. Delta U = Q Vertical line on P-V graph

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Isovolumetric process

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Can occur without additional energy. Many of them have high Ea and therefore go slowly and may not go to completion.

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Spontaneous processes

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Require energy to happen. Often coupled to spontaneous reactions so they go

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Nonspontaneous processes

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...

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*7.2 States and State Functions*

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Describe system in an equilibrium state. Useful for comparing one equilibrium state to another

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State functions

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Describe quantitatively the pathway taken from one pathway to another

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Process functions

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Pressure, density, temperature, volume, enthalpy, internal energy, Gibbs free energy, entropy

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Examples of state functions

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Work and heat

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Examples of process functions

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25 Celsius, 1 atm pressure, 1M For consistency in kinetics, equilibrium, and thermochem

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Standard conditions for Thermochemistry and why are they used?

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Most stable form of a substance under standard conditions e.g. H2(g), H2O(l), NaCl(s)

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Standard state of a substance

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Delta H naught Delta S naught Delta G naught

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Changes in enthalpy, entropy and free energy under standard conditions

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Show standard and nonstandard states of matter for a given substance in an isolated system, determined by temperature and pressure.

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Phase diagrams

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Reversible, and equilibrium of phases will be reached at any temp and pressure. e.g. some ice absorbs heat and melts. This heat is removed from the surrounding water, and an equal amount freezes. Similar thing happens liquid and gas phases of water in a closed container.

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Phase changes

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Temperature related to average kinetic energy of the molecules that make up the substance. Therefore molecules have a range of instantaneous KE values.

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Temperature and KE

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Some molecules at surface of liquids gain enough KE to escape into gas phase.

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Evaporation/vaporization

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Specific type of vaporization that occurs above the BP and involves vaporization through the entire volume

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Boiling

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Gas goes back into liquid phase. Occurs due to lower temp or higher pressure. In covered or closed container, escaping molecules trapped, exert a countering pressure and they go back to liquid

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Condensation

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The pressure that the gas exerts over the liquid at equilibrium. Increased with temp increased because KE increases and more molecules enter gas phase

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Vapor pressure of the liquid

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When vapor pressure of the liquid equals the ambient pressure

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Boiling point

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The number of motions about an equilibrium position that a atom/molecule can do, which increases as temperature of solid increases.

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Microstates

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Pure crystalline substances have precise MPs Amorphous solids (glass, plastic, wax, chocolate) melt or solidify over a large range of temperatures.

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MP of pure vs. amorphous solids

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Device used to purify a product that is heated under reduced pressure, causing it to sublime (go to gas) Desired product usually more volatile than impurities so gas is product. Gas deposits onto cold finger, and cold water cools the gas, yielding the pure solid

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Cold finger

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Indicate phase boundaries, showing the temp and pressure where equilibrium between phases happens.

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Phase diagrams - lines of equilibrium

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Where the three phase boundaries meet. At this temperature and pressure, all three phases in equilibrium

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Triple point

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Where the liquid-gas line ends. At and beyond this point, no distinction between liquid and gas (supercritical fluids)

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Critical point

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As liquid is heated in closed system, density of liquid decreases and density of gas increases. At critical point, two densities are equal and there is no distinction between the two phases.

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How do supercritical fluids work?

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Related to the average KE of the particles of a substance.

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Temperature (T)

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Related to average KE of the praticles as well. Also depends on how much of the substance is present. If substance's thermal energy increases, temperature also increases

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Thermal energy (enthalpy)

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Transfer of energy from one substance to another because of a difference in temperature.

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Heat (Q)

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A law that if two systems are separately found to be in thermal equilibrium with a third system, the first two systems are in thermal equilibrium with each other; that is, all three systems are at the same temperature. Also known as thermodynamic equilibrium.

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Zeroth law of thermodynamics

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Delta Q > 0

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Endothermic

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Delta Q < 0

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Exothermic

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Joule or calorie (1 cal = 4.184 J)

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Unit of heat

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Constant pressure. Assume this.

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When is enthalpy (delta H) equivalent to heat?

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Process of measuring transferred heat.

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Calorimetry

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q = mc(delta T) where c is the specific heat and T is the change in temp (Celsius or Kelvin)

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Heat absorbed or released equation

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Amount of energy required to raise the temperature of one gram of the substance by one degree Celsius (or kelvin). c(h2o) = 1 cal/g*K

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Specific heat

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mass times specific heat (= mc)

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Heat capacity

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Insulated container covered with lid. Contains solution where some reaction or physical process is occurring. Atmospheric pressure stays constant and temperature can be measured as the reaction progresses.

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Constant-pressure calorimeter (coffee-cup)

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Sample of matter (usually hydrocarbon) goes in steel decomposition vessel, vessel filled with O2. Vessel placed in insulated container with known mass of water. Contents of decomposition vessel ignited, and heat evolves --> heat of combustion reaction. W = P(delta V), and since isovolumetric process, no work done. System = sample + O2 + steel vessel. Surroundings = water

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Constant-volume calorimetry (bomb calorimeter)

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No heat exchanged between calorimeter and rest of universe, so Q = 0 Delta Usys + Delta U surroundings = Delta U calorimeter = Qcal - Wcal = 0 Delta Usys = -DeltaUsurr No work done, so qsys = -qsurroundings (m)steel(c)steel*deltaT + m(oxygen)c(oxygen)*deltaT = -m(water)c(water)*deltaT

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Math for bomb calorimeter

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No heat exchange between calorimeter and universe (therefore adiabatic process), but there is exchange between decomposition vessel and water.

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Heat exchange in bomb calorimeter

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use q = mc(delta T) and the fact that qcold = -qhot. Solve for Tf

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Question type - two substances, equilibrium temperature?

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Temp rises until MP or BP reached. Temp constant as compound converted to next phase. After conversion, temp begins to rise again. Temp constant during phase change because substance absorbs energy to allow particles to overcome attractive forces to go into the next phase.

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Heating curve

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Solid to liquid - heat of fusion Liquid to gas - heat of vaporization

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Solid-liquid and liquid-gas enthalpy values for phase changes

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q = mL L is the latent heat (enthalpy of an isothermal process - cal/g)

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Find heat used in a phase change

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sum of heat for changing temp of each respective phase, and heats associated with phase changes.

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Total amount of heat to cross multiple phase boundaries

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= H. A state function. delta H of a system can be calculated by comparing the enthalpy of the final state to the enthalpy of the initial

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Enthalpy

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= Hproducts - Hreactants Positive delta Hrxn = endothermic Negative delta Hrxn = exothermic

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delta H equation

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Enthalpy change accompanying rxn being carried out under standard conditions Delta H naught rxn = sum(std heat of formation, products) - sum(std heat of formation) reactants

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Standard heat of formation (explain, and formula for rxn)

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Enthalpy changes are additive. Note: Enthalpy change for the reverse of any reaction has same but opposite magnitude. Hess's law works for ANY STATE FUNCTION INCLUDING ENTROPY AND GIBBS FREE ENERGY

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Hess's law

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Average energy required to break a particular type of bond between atoms *in the gas phase* Units of kJ/mol of bonds broken

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Bond enthalpy (bond dissociation energies)

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Opposite of bond breaking, same magnitude of energy, but negative (energy released when bonds formed). delta H naught rxn = sum(deltaH bonds broken) - sum(delta H bonds formed)

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Bond formation

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Add up the delta H for bonds breaking (positive) and bonds forming (negative). Multiply value for the # of bonds broken or formed. p. 224. Look at it now.

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Questions with finding enthalpy

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delta H naught comb = enthalpy change associated with combustion of fuel. Measurements require reaction to be spontaneous and fast - combustion reactions are great for this.

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Standard heat of combustion

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Energy is going from being localized/concentrated to being spread out or dispersed.

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Entropy - idea

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Energy will spontaneously disperse from being localized to becoming spread out if it is not hindered from doing so.

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Second law of thermodynamics

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Measure of spontaneous dispersal of energy at a specific temperature Where Q = Qrev (heat gained or lost in a reversible process) Units usually J/mol*K

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Entropy and formula

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delta Suniverse = delta S sys + delta S surr > 0

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Entropy of sys and universe

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It is a state function, therefore can calculate delta S naught reaction. delta S naught rxn = sum(delta S naught formation products) - sum(delta S naught formation reactants)

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Entropy as a state function

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Measure of the change in enthalpy and the change in entropy as a system undergoes a process. Indicates if reaction is spontaneous or nonspontaneous. delta G = delta H - T delta S T is temperature in Kelvin T(delta S) represents amount of energy absorbed by system when entropy increases reversibly.

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Gibbs free energy

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delta G negative = spontaneous delta G positive = nonspontaneous delta G zero = equilibrium where delta H = T(delta S)

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Gibbs free energy sign and spontaneity

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For example, for equilibrium between gas and solid delta G = G(g) - G(s) = 0 *Always has to be zero*!@#*()!*@()

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Gibbs free energy change for phase equilibria

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Effect on delta G = delta H - T (delta S) Spontaneous --> + deltaH + deltaS (at high T); - deltaH + deltaS (all temp); - deltaH - deltaS (at low T) Nonspontaneous --> + deltaH - deltaS (all temp)

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Effects of delta H and delta S on spontaneity

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This is an endothermic process - need to put energy in. delta H and delta S will be positive. Therefore reaction will only be spontaneous (negative delta G) if T(deltaS) is larger than delta H. This only occurs at temperatures above 100 degrees. At 100 degrees, delta H - T(deltaS) = 0, and equilibrium established between liquid and gas phases - water's vapor pressure equals ambient pressure

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Example of water boiling.

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delta G determines spontaneous or not. Ea determines reaction rate.

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Ea vs. delta G

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Under standard state conditions. deltaG naught formation, is the free energy change occurring when 1 mole of a compound in its standard state is made from its respective elements in their standard states under standard state conditions.

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Standard free energy

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= 0

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Standard free energy of formation for any element under standard state conditions

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= -R T ln(Keq) R = ideal gas constant, T temperature in kelvins. Higher Keq = more positive ln = more negative standard free energy change = more spontaneous reaction

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Standard free energy change formula

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They no longer apply

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Standard state conditions AFTER REACTION BEGINS

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delta G (rxn) = (standard free energy change)rxn + RTlnQ = RTln(Q/Keq) If Q/Keq is less than one (Q Keq), then natural log is positive and free enegy change is positive. Reaction will go in the reverse direction until equilibrium is reached. If ratio = 1, then free energy change is 0 (ln 1 = 0)

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Free energy change for a reaction IN PROCESS

Physical Chemistry Chapter 7 (Thermochemistry) – Flashcards

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