Exam 1 Review Flashcards
Enthalpy
A thermodynamic quantity = total heat content of a system.
- *H = qp**
- enthalpy = heat at constant pressure*
H = U + PV *enthalpy = internal energy + (pressure)(volume)*
*State function
*Reminder: U = q + w
internal energy = heat + work
Entropy
Represented by S
A measure of energy dispersal (disorder
Three Laws of Thermodynamics
- Energy is not lost or gained during chemical reactions, it only changes form.
- Entropy (S; free energy) of the universe is continually increasing.
- Entropy of a perfect crystal at absolute zero is zero (used for standard entropy, Sº).
Is the ΔSsys positive or negative when…?
a) a solid melts
b) ions combine to form a solid
c) a gas forms during a reaction
d) four reactants combine to make two products
a) a solid melts → +ΔSsys
b) ions combine to form a solid → -ΔSsys
c) a gas forms during a reaction → +ΔSsys
d) four reactants combine to make two products → -ΔSsys
Spontaneous reactions cause ΔSuniv to increase or decrease?
Spontaneous reactions → +ΔSuniv
Gibbs free energy change
ΔG = ΔH -TΔS
Thermodynamic function that expresses sponaneity of a process more directly.
At constant T & P:
ΔG < 0Spontaneous in forward direction.
ΔG = 0 System is at equilibrium.
ΔG > 0 Nonspontaneous in forward direction.
Predicting sign of ΔG using signs of
ΔH (enthalpy) and ΔS (entropy)
ΔG = ΔH - TΔS
-ΔG = -ΔH - TΔS
Always spontaneous.
<em>(Think: Always <strong>spontaneous</strong> when <strong>only</strong> <strong>S</strong> is <strong>positive</strong>.)</em>
+ΔG = ΔH - T(-ΔS)
Always nonspontaneous.
?ΔG = -ΔH - T(-ΔS)
-ΔG = TΔS < ΔH & Low T = spontaneous
+ΔG = TΔS > ΔH & High T = nonspontaneous
<em>(Think: Both “<strong>low</strong>” (negative) and <strong>low</strong> T = <strong>spontaneous</strong>.)</em>
?ΔG = ΔH - T(ΔS)
+ΔG = TΔS < ΔH & Low T = nonspontaneous
-ΔG = TΔS < ΔH & High T = spontaneous
(Think: Both “<strong>high</strong>” (positive) and <strong>high</strong> T = <strong>spontaneous</strong>.)
Differences between gases, liquids, and solids
Disorder:
Gases = total disorder
Liquids = some disorder
Solids = ordered arrangement
Motion:
Gases = total freedom of motion
Liquids = some freedom of motion, clusters
Solids = fixed positions, little/no motion
Spacing:
Gases = far apart
Liquids = close together
Solids = packed close together
Intermolecular attractions:
Gases = none
Liquids = strong
Solids = very strong
Kinetic energy:
Gases = high
Liquids = between high/low
Solids = low
Energy and phase change
q = m*ΔT*s
used when there is a temperature change
q = ΔHphase change * n
used when there is a phase change/no temperature change
Differences between hydrogen bonds, dipole-dipole interactions, and London forces
Hydrogen bonds:
H and NOF
Proton attracted to highly electronegative atom
Dipole-dipole:
Polar molecules are attracted to each other by oppositely charged ends
London forces:
Temporary dipoles attract each other, very weak and short-lived interaction
Why do molecules with H-bonds have higher boiling points?
Hydrogen bonds are stronger than other intermolecular interactions
*NOTE: Larger, more complex molecules have greater London forces, causing a difference in boiling points between H-bonded molecules
Simple-, body-, and face-centered unit cells
Simple:
Coordination # = 6
# of atoms = 1
Body:
Coordination # = 8
# of atoms = 2
Face:
Coordination # = 4
# of atoms = 12
Ionic, covalent, and molecular solids
Ionic:
Lattice points = like ions
Overlapping unit cells of cations/anions
(e.g. NaCl)
Covalent:
Lattice points = atoms
Connected by covalent bonds
(e.g. graphite/C)
Molecular:
Lattice points = molecules
Connected by intermolecular forces
(e.g. ice)
“Like dissolves like”
Chemicals with similar polarities will dissolve in each other
Enthalpy of a solution
Heat change when a solute is dissolved in a solvent
