Kaplan MCAT General Chemistry Chapter 7

Thermochemistry

The study of energy changes that occur during chemical reactions and changes in state

System

The matter that is being observed.

The total amount of reactants and products in a chemical reaction.

Surroundings or environment

Everything outside of the system

Isolated system

The system cannot exchange energy (heat or work) or matter with the surroundings; for example, and insulated bomb calorimeter.

isolated—- as in by yourself— as in with no one and nothing— so this one cannot exchange both

Open system

The system can exchange energy (heat and work) and matter with the surroundings

Closed system

The system can exchange energy (heat and work) but not matter with the surroundings

Process (system)

A change in one or more of its properties (such as concentrations of reactants or products, temperature, or pressure).

First law of thermodynamics and related equation

Energy can be transferred and transformed, but it cannot be created or destroyed.

U-internal energy of the system

Q-heat added to system

W-work done by system

Isothermal process

Constant temp. implies that the total internal energy of the system (U) is constant

no change in temp.; change in ΔU=0; Q=W (heat added to the system equals work done by the system)

Adiabatic process

A process in which no heat is transferred to or from the system by its surroundings

When Q = 0, the first law simplifies to ΔU = -W (the change in internal energy of the system is equal to work done on the system)

Isobaric process

Occur when pressure of the system is constant. Does not alter the first law of thermodynamics.

Isochoric (isovolumetric)

A process in which there is no change in volume.

W=0, ΔU=Q

no work done

Spontaneous process

One that can occur without having to be driven by energy from outside source.

coupling

A common method for supplying energy for nonspontaneous reactions is by ___________ nonspontaneous reactions to spontaneous ones

State functions

Describe the system in an equilibrium state.

Path independent

When I am under Pressure and feeling Dense, all I want to do is watch TV and get HUGS.

Pressure (P), density (p), temperature (T), volume (V), enthalpy (H), internal energy (U), Gibbs free energy (G), and entropy (S).

Entropy Equation

What are standard conditions? and what should it not be confused with

The standard conditions defined for measuring the enthalpy, entropy, and Gibbs free energy of a reaction.

25 C (298 K), 1 atm pressure, and 1 M concentration.

don't confuse this with STP which is used in gas law calculations: at STP temp is 273k and pressure is 1atm

Standard state of a substance

The most stable state a substance can exist in at standard conditions.

For example, H2 (g), H20 (l), O2 (g), NaCl (s), C(s)

Phase diagrams

• draw and label each phase and each point of importance

Graphs the phase equilibria as a function of temperature and pressure for a system.

Phase changes (types)

Exist as characteristic temperature and pressures (they are reversible):

1) Fusion (melting) and freezing (crystallization and solidification)

2) Vaporization (evaporation or boiling) and condensation

3) Sublimation or deposition occur at the boundary between the solid and gas phase.

evaporation (vaporization)

A transformation from a liquid to a gas

-ex: boiling (happens when a liquid is heated to become a gas at its boiling point)

condensation

Gas to liquid

- lower temperature and higher pressure

melting (fusion)

solid to liquid

freezing (crystallization or solidification)

liquid to solid

sublimation

solid to gas

depositon

gas to solid

Critical point

• what to note about it

• what is it called here

• draw phase diagram labeling x and y axis

The temperature and pressure at which the gas and liquid states of a substance become identical and form one phase

• supercritical fluid

- heat of vaporization at this point and all temps and pressures above this point is zero

Triple point

The point where all three phases of matter exist in equilibrium.

Heat

The transfer of energy that results from differences of temperature between two substances.

Temperature

The scaled measure of the average kinetic energy of a substance. (how hot or cold something is)

Heat vs Temperature

Heat is a specific form of energy that can enter or leave a system. (process function)

Temperature is a measure of the kinetic energy of the particles in a system.

Zeroth law of thermodynamics- 2 definitions

objects are in thermal equilibrium only when their temperatures are equal

if two thermodynamic systems are both in thermal equilibrium with a third system, then the two systems are in thermal equilibrium with each other

endothermic

Absorbs heat (ΔQ>0)

exothermic

Releases heat (ΔQ<0)

what is the unit of heat

Enthalpy (ΔH)

what happens when substances of different temperatures are brought into thermal contact- meaning that there is a physical arrangement allowing for heat transfer

unit of energy called the joule. one joule us equal to 4.184 cal

The heat content, Q, of a system at constant pressure

energy will move from warmer to cooler substance

Calorimetry

The precise measurement transferred heat

constant pressure calorimeter

Insulated container covered with a lid and filled up with a solution in which a reaction or some physical process (such as dissolving) is occurring.

Equation of heat transfer-no phase change

q=mcAT, where m=mass, c=specific heat, T = change in temp.

Specific heat and heat capacity

• what does heat capacity mean

the amount of energy required to raise the temperature of one gram of a substance by one degree Celsius (or one degree kelvin).

heat capacity=mc

bomb calorimeter

a device for measuring the heat evolved in the combustion of a substance under constant-volume conditions

constant volume calorimetry

the measurement of heat at a constant volume to find the change in internal energy of the system; requires the use of a "bomb" calorimeter

Draw a heating curve graph

• make note about phase change regions

change regions do not undergo changes in temp— at constant temp

ΔHfus

enthalpy of fusion

- must be used when transitioning at the solid-liquid boundary

ΔHvap

enthalpy of vaporization

-must be used when transitioning at the liquid-gas boundary

Calculating heat transfer (q) during phase changes?

Energy is still required but the change in temperature is 0.

q= internal energy

L = latent heat (enthalpy of isothermal process)

m = mass

cannot use mcdeltaT equation because temperature is constant during a phase chnage— would be erroneous to assume q would be zero

enthalpy

The heat content of a system at constant pressure

Enthalpy change of reaction, ΔHrxn for Endo thermic and exothermic

+ ΔH = endothermic

-ΔH= exothermic

Hess's Law

Standard enthalpy of formation— whats the symbol and definition

enthalpy required to produce one mole of a compound from its elements in their standard states

by definition, ΔHf of an element in its standard state is 0

standard heat of a reaction (ΔHrxn) definition and equation

The change in enthalpy of a reaction under standard conditions

Hess's law

• what is it and what does it apply to (practice equations for Hess’s law— how to find overall heat of reaction given multiple equations)

states that enthalpy changes of rxn are additive—- applies to any state function including entropy and Gibb’s free energy

When doing problems, make sure to switch signs when you reverse the equation

Bond Enthalpy Equation— once again practice using these equations

Energy absorbed-energy released

Bond dissociation energy

• what is important to note and use mnenomic device

• BARF- broken absorbed, released formed in relation to bond energies

The average energy required to break a particular type of bond between atoms in a gas phase.

formation exothermic, energy release (-)

dissociation endothermic

standard heat of combustion— once again, practice using this in equation

ΔH°(comb) is the enthalpy change associated with the combustion of a fuel

most combustion appears in presence of O2 in atmosphere

2nd law of thermodynamics

energy spontaneously disperses from being localized to becoming spread out if it is not hindered

Entropy

• definition and equation

a measure of the disorder of a system, measures at a specific temperature

qrev/T where Qrev is the heat that is gained or lost in a reversible process

entropy of the universe

• meaning and equation

entropy of the universe is always increasing

standard entropy change

use Hess’s law

entropy in phases in order from least to most

the amount of entropy increases from solid--> liquid---> gas

Change in Gibbs free energy

Goldfish are Horrible (minus sign) without tarter sauce

gibbs free energy chart

ΔG is temperature dependent when ΔH and ΔS have the same sign

remember order like HS G— his girlfriend

when the signs are opposite, the third and first sign match, example” -+-

Standard Gibbs free E from eq const

the greater the value of Keq, the more negative the standard free energy change (the more spontaneous the reaction)

exergonic reaction

A spontaneous chemical reaction in which there is a net release of free energy.

endergonic reaction

Reaction that absorbs free energy from its surroundings. (nonspontaneous)

endergonic/exergonic

endothermic/exothermic

endergonic/exergonic - Gibbs free energy

endothermic/exothermic - enthalpy

Standard Gibbs free energy of reaction

Standard Enthalpy of Reaction

Gibbs free energy from reaction quoteint

Generalized Enthalpy of Reaction

the _______ the alkane the more numerous the combustion product

the __larger__ the alkane the more numerous the combustion product

thermodynamics vs. kinetics

thermodynamics and kinetics are separate. when a reaction is thermodynamically spontaneous, it has no bearing on how fast it goes. it only means that it will proceed eventually