UT Austin Chem 1 Thermodynamics

Open system

Matter and energy can be exchanged

Closed system

Energy can be exchanged, but matter cannot

Isolated System

Matter and energy cannot be exchanged

Temperature, density, pressure

Intensive properties - define the system and remain the same regardless of sample size

Volume, mass, moles, energy

Extensive properties - depend on the amount of substance

State functions

depend only on the starting and final values of the system: include delta H, delta G, delta U, and delta S

Path functions

follows the path of the process, such as w (work), and q (heat).

Internal Energy

U

If delta U is positive…

The system gains energy

If delta U is negative…

The system loses energy (flows into surroundings)

First law of Thermodynamics

Energy cannot be created or destroyed, therefore the change in internal energy of the system (delta U) will be equal to the negative/positive counterpart of delta U surroundings

Delta U = q + w

Change in internal energy equals heat plus work

Work

the energy transferred as a result of a force acting through a distance

equations for work

W= -P(delta V) ; W=-(delta n of gas molecules) RT

Q is positive

heat is added to the system

Q is negative

heat is removed from the system

W is positive

work is done on the system

W is negative

work is done by the system (onto the surroundings)

Delta G is positive

Non-spontaneous and endothermic

Delta G is negative

spontaneous and exothermic

Delta S is positive (entropy)

increasing disorder

Delta S is negative (entropy)

increasing order

Delta H is negative

Exothermic

Delta H is positive

endothermic

What does H represent

H is enthalpy - the total heat content in the system

Equation for Delta H

Delta H = Delta U + P(delta V)

In a closed system…

Volume is constant, so Delta U = q(v)

In an open system…

pressure is constant, so Delta H = q(p)

If a gas expands

delta V is going to be positive as work is done on the surroundings, so w is negative

If a gas is compressed

Delta V is negative as work is done on the system , so w is positive

101.3 J

equals 1 L *atm

R value for thermodynamics

8.314 j/mol k

Heat Capacity (C)

The amount of heat required to to raise the temperature of a substance by 1 degree Celsius 

Specific Heat Capacity

The amount of heat required to raise 1g of a substance by 1 degree Celsius 

Molar Heat Capacity

The heat required to raise the temperature of 1 mole of a substance by 1 degree Celsius

Heat capacity equation in grams

q= mc(delta T)

Heat capacity equation in moles

q= nCm(deltaT)

Factors that influence heat capacity

C increases with size, amount, and IMF

If q1=q2

then mcat1=mcat2 or ncmt1=ncmt2 (often used in calorimetry problems)

Calorimetry

the technique we se for measuring the heat involved in a chemical or physical change

Constant Pressure Calorimeter

Also called Coffee Cup calorimeter- an insulated container covered with a lid and filled with.a solution in which a reaction or physical process is occurring. System is open to the atmosphere. Delta H= Qp. Helps determine spontaneity of a reaction based on delta H.

Delta H on a graph

is the difference between the start and end of the line, exothermic reactions will have a negative H as ending H will be smaller than initial H.

Constant Volume Calorimeter

Also called Bomb Calorimeter - allows us to measure the internal energy change of a reaction. Reaction is done in a bomb under a substance such as water. Delta U is equal to Qv

To calculate q for calorimeters:

q= mc(delta T) + c(delta T) → adds heat of liquid surrounding the reaction and the heat or the calorimeter

Equation for change in enthalpy

delta H= delta U + P (delta V)

Hearing curve

The temperature of a substance increases until it reaches its melting/boiling point, then the temperature will remain constant until the substance reaches the next phase

Solid warming up

Q= mC(solid)(delta T)

Solid melting to liquid

q=mass times delta H of fusion

Liquid warming up

q=mC(liquid)(delta T)

Liquid vaporizing to gas

Q=m times delta H of vaporization

Gas warming up

Q=mC(gas)(delta T)

The four methods for determining the change in enthalpy of a reaction are…

1 doing the experiment in a lab 2 using bond enthalpies 3 using standard standard enthalpies of formation 4 using Hess’ law to sum enthalpies that lead to the desired reaction

Bond enthalpies

energy needed to break one mol do fate bond to give separated atoms - everything being in the gas state - breaking bonds is an endothermic reaction - DeltaH (rxn)= the sum of the bond enthalpies of the reactants minus the sum of the bond enthalpies of the products → result is an approximation

Heat of formation

enthalpy change that occurs when one mole of a compound is formed from its elements in their most stable forms (standard rates) under standard conditions of 298k and 1atm pressure. Delta H (rxn)= the sum of the heat of formation (times moles) of the products minus the sum of the heat of formation (times moles) of the reactants.

Hess’s Law

states that because enthalpy is a state function, if two or more chemical reactions can be added together to give a specific reaction, their change in enthalpies can be added together give the change in enthalpy for that specific reaction. 

Rules for Heat of Formation Reactions

1)They produce 1 mole of substance, 2) reactants are elements in their standard states 3) under standard conditions 1atm and 298K. Heat of formation is more. Accurate than bond energies because it accounts for all phases.

Standard states for 7 diatomic substances

H2,N2,O2,F2, Cl2 = gasses ; Br2=liquid ; I2 = solid

elements in their standard states…

have a change in enthalpy of formation of 0

Entropy (S)

The measure of the spontaneous dispersal of energy at a specific temperature- related to the number of microstates available to a system

Gibbs Free Energy (determines spontaneity)

Delta G = delta H - T delta S

Boltzmann Equation

S=klnW

As microstates increase…

Entropy increases (increase in randomness)

As volume or size of the molecules increases…

Increases the number of microstates

Second Law of Thermodynamics

In any spontaneous reaction, the total entropy of the universe increases

Delta S (universe)= delta S (system)+ delta S(surroundings)

How to find change in entropy

Delta S = q/T ; Delta S (surroundings)= -Delta H(system)/T

If the sign of delta S (surroundings) is negative…

It is an endothermic process (loses randomness, becomes more solid, etc)

As a substance goes from solid →liquid→ gas

Delta S (system) is positive

As molar weight increases…

Entropy generally increases

Mixtures are more disordered…

Than pure substances

A floppy structure (such as hydrocarbons) has more entropy than…

Rigid structures (such as ring structures)

If the products of a reaction has less moles (gas) than the reactants…

Then there is less entropy and S is negative

Change in entropy in a reaction (same as enthalpy

Delta S(reaction) = sum of n times delta S products - sum of n times delta S reactants

Change in Gibbs Free energy equation

Delta G (rxn) = total enthalpy of reaction - T times total entropy of reaction

If delta G is negative…

If delta G is positive…

If delta G is zero…

The reaction is spontaneous

The reaction is non spontaneous

The system is in equilibrium

Entropy and enthalpy are positive

The reaction is spontaneous at high temperatures

Enthalpy is positive and entropy is negative

Non spontaneous at all temperatures

Enthalpy is negative and entropy is positive

Spontaneous at all temperatures

Enthalpy and entropy are negative

Spontaneous at low temperatures

To find the transition temperature

Set delta G to 0 and solve for temperature→ if both h and s are positive, then the reaction is spontaneous above the transition temperature and vice versa

Gibbs Free Energy of Formation

Defined as the change in GFE when one mole of a compound is formed from its elements in their standard states. Negative G (formation) means the reaction is spontaneous and the more negative G the more stable the substance