MCAT Study Guide Chemistry Ch. 6 – System and Surroundings 2017-08-15T06:45:06+00:00

I.          6.1:  SYSTEM AND SURROUNDINGS

A.     THE ZEROTH LAW OF THERMODYNAMICS

1.     Establishes definition of thermal ⇋:  when systems are in thermal ⇋ with each other, their temps must be the same

2.     When bodies of different T are brought into contact with one another, heat will flow from the body with the ↑T to the body with the ↓T

B.     THE FIRST LAW OF THERMODYNAMICS

1.     The total energy of the universe is constant (conservation of energy)

2.     Adiabatic – refers to a process that occurs without the transfer of heat or matter

 

II.         6.2:  ENTHALPY

A.     Enthalpy (H)

Measure of the heat energy that is released or absorbed when bonds are broken and formed during a reaction

1.     When bonds are formed, energy is released (ΔH < 0)

2.     When bonds are broken, energy is needed (ΔH > 0)

B.     ΔH = heat of reaction = Hproducts – Hreactants

C.    H = U + PV

D.    Enthalpy is affected by phase change, intermolercular forces, strength of intramolecular forces

 

III.          6.3:  CALCULATION OF ΔHRXN

A.     STANDARD CONDITIONS

1.     T = 25° C, P = 1 atm, concentration = 1 M

2.     Don’t confuse with STP, where T = 0° C!!

B.     HEAT OF FORMATION

1.     Standard heat of formation (ΔH°f) – amount of energy required to make one mole of a compound from its constituent elements in their natural or standard state at 1 atm (no standard temperature)

a)     ΔH°f = 0 for elements in their standard state (like O2, Cl2)

b)     ΔH°f for O = 249 kJ/mol, because it takes energy to break the O=O

c)     ΔH°f for water is exothermic

C.    HESS’S LAW OF HEAT SUMMATION

1.     Hess’s Law – if a reaction occurs in several steps, then the sum of energies absorbed or given off in all the steps will be the same as that for he overall reaction

a)     If the reaction is reversed, the sign of ΔH is reversed

b)     If an equation is multiplied by a coefficient, then ΔH must be multiplied by that same value

D.    SUMMATION OF AVERAGE BOND ENTHALPIES

1.     Bonds have characteristic enthalpies (BDE), and the sum bond energies equals the total ΔH

2.     ΔHrxn = Σ(BDE broken) – Σ(BDE bonds formed)

 

IV.          6.4:  ENTROPY

A.     THE SECOND LAW OF THERMODYNAMICS

1.     All processes tend to run in a direction that leads to maximum disorder (entropy)

2.     Entropy (S) is predictable:

a)     Gases have more entropy than liquids which have more entropy than solids

b)     Particles in solution have more entropy than undissolved liquids

c)     Two moles of a substance have more entropy than one mole

d)     ΔS for a (-) reaction has the same magnitude but opposite signs as the (+) reaction

3.     ΔS = q/T

B.     THE THIRD LAW OF THERMODYNAMICS

1.     Absolute zero = zero entropy

 

V.          6.5:  GIBBS FREE ENERGY

A.     Gibbs free energy (ΔG)

Energy available to do useful work from a chemical reaction; determines spontaneity of a reaction

1.     ΔG = ΔH – TΔS

a)     ΔG < 0  → spontaneous reaction

b)     ΔG = 0  → reaction at ⇋

c)     ΔG > 0  → nonspontaneous reaction

B.     ΔG AND T

1.     Entropy depends directly on temperature

a)     At low T, entropy does not have much influence on free energy and ΔH is dominant factor in determining spontaneity

b)     As T ↑, S becomes more significant relative to ΔH in determining spontaneity

2.     In general, the universe tends toward ↑ disorder and ↑ stable bonds (negative ΔH)

 

VI.          6.6:  REACTION ENERGY DIAGRAM

A.     KINETICS VS THERMODYNAMICS

1.     Thermodynamics predicts spotaneity of reaction

2.     Kinetics predicts rate of reaction

B.     REVERSIBILITY

1.     Reaction diagrams can be drawn by using mirror image of the reverse reaction

2.     Note that Ea is different for each one

MCAT Study Guide Chemistry - Kim Matsumoto


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