Class 2: Introduction to Free Energy Flashcards
Connect your understanding of potential energy from Chem 150 to the new concept of bond dissociation enthalpy.
Potential Energy Review:
- Potential energy is the energy stored within a system due to position/configuration
- Molecules/bonds have potential energy based on relative atomic positions
- Breaking bonds increases potential energy (endothermic)
Bond Dissociation Enthalpy:
- The energy required to break a specific covalent bond
- Quantifies the bond strength and potential energy stored
- Higher dissociation enthalpy = stronger bond, more energy to break
Connections:
- Bond dissociation enthalpy is a specific type of potential energy
- It represents the potential energy stored within a covalent bond
- Breaking a bond releases this stored potential energy
- The stronger the bond, the higher its dissociation enthalpy/potential energy
Energy Changes:
- Energy must be supplied to break bonds (endothermic)
- This energy overcomes the potential energy barrier of the bond
- Once broken, the atoms have higher potential energy in separated state
- Bond formation is an exothermic release of this potential energy
Key Points:
- Bond dissociation enthalpy quantifies potential energy of a covalent bond
- It’s the energy input required to overcome the bond’s potential energy barrier
- Stronger bonds have higher dissociation enthalpies/potential energies
- Bond breaking increases potential energy; bond forming releases it
Let me know if any part needs further clarification!
Calculate the enthalpy of a reaction (ΔH°) from bonds broken and bonds formed.
Given:
- Bond dissociation enthalpies for all bonds broken in reactants
- Bond dissociation enthalpies for all bonds formed in products
Step 1: Sum the bond dissociation enthalpies for bonds broken
- This gives the total energy input to break reactant bonds
Step 2: Sum the bond dissociation enthalpies for bonds formed
- This gives the total energy released from forming product bonds
Step 3: Calculate ΔH°
- ΔH° = ΣBond Dissociation Enthalpies (Bonds Broken)
- ΣBond Dissociation Enthalpies (Bonds Formed)
Notes:
- Energies for breaking bonds are positive (endothermic)
- Energies for forming bonds are negative (exothermic)
- ΔH° positive means overall endothermic
- ΔH° negative means overall exothermic
- Units of ΔH° are kJ/mol
Example:
CH4 + 2O2 → CO2 + 2H2O
Bonds Broken: 4(C-H), 2(O=O)
Bonds Formed: 2(C=O), 4(O-H)
ΔH° = [4(413) + 2(498)] - [2(799) + 4(463)]
= 1652 + 996 - 1598 - 1852
= -802 kJ/mol
Describe the difference between endothermic and exothermic reactions.
Endothermic Reactions:
- Absorb heat energy from the surroundings
- Temperature of the surroundings decreases
- Require input of energy to proceed
- Have a positive enthalpy change (ΔH > 0)
- Examples: Photosynthesis, melting, cooking
Exothermic Reactions:
- Release heat energy to the surroundings
- Temperature of the surroundings increases
- Release energy as they proceed
- Have a negative enthalpy change (ΔH < 0)
- Examples: Combustion, neutralization, respiration
Key Differences:
Endothermic | Exothermic
Absorbs Heat | Releases Heat
ΔH is positive | ΔH is negative
Energy is absorbed | Energy is released
Surroundings cool | Surroundings warm
Requires energy input | Releases energy
Energy Flow:
- Endothermic: Energy flows from surroundings → system
- Exothermic: Energy flows from system → surroundings
Overall:
- Endothermic reactions take in energy to break bonds
- Exothermic reactions release energy from forming bonds
- Enthalpy change determines endo/exothermic nature
Define enthalpy and explain how it relates to free energy.
Enthalpy (H):
- A state function that combines internal energy (U) and pressure-volume work (PV)
- H = U + PV
- Useful for measuring energy changes in reactions at constant pressure
Free Energy (G):
- The energy available to do non-expansion work
- Accounts for enthalpy (H) and entropy (S)
- G = H - TS (T is absolute temperature)
Relationship between H and G:
- Enthalpy term (H) represents total heat energy of the system
- Entropy term (-TS) accounts for molecular disorder/randomness
- Free energy (G) is the usable portion of the enthalpy after entropy effects
Free Energy Changes:
- ΔG < 0 means a spontaneous, energy-releasing process
- ΔG > 0 means a non-spontaneous, energy-absorbing process
- ΔG is determined by both enthalpy (ΔH) and entropy (ΔS) changes
Significance:
- Reactions are driven by decreasing free energy (ΔG < 0)
- More negative ΔG means greater driving force
- Enthalpy and entropy changes contribute in opposite ways to ΔG
Key Point: Free energy relates the total enthalpy of a system to its randomness/disorder, determining if energy is available for useful work.
Describe the difference between spontaneous and non-spontaneous processes.
Spontaneous Processes:
- Occur naturally on their own without outside assistance
- Have a negative Gibbs free energy change (ΔG < 0)
- Are thermodynamically favored
- Examples: Bomb explosions, combustion reactions, ice melting
Non-Spontaneous Processes:
- Require continual input of energy to occur
- Have a positive Gibbs free energy change (ΔG > 0)
- Are thermodynamically not favored
- Examples: Electrolysis, photosynthesis, charging a battery
Key Differences:
Spontaneous | Non-Spontaneous
ΔG < 0 | ΔG > 0
Favored by nature | Not favored
Releases free energy | Absorbs free energy
Occurs naturally | Requires energy input
Increases disorder | Decreases disorder
Driving Forces:
- Spontaneous: Driven by tendency to minimize free energy
- Non-Spontaneous: Driven by increasing free energy
Effects of ΔH and ΔS:
- For spontaneity, ΔG must be < 0
- ΔG = ΔH - TΔS
- Spontaneous if ΔH is negative OR if TΔS is large positive
The sign of ΔG determines if a process is spontaneous or non-spontaneous at that condition.
Define what a state function is and give examples.
State Function:
- A property that depends only on the present state of the system
- Not dependent on how that state was reached or the path taken
- Has a precise value that can be calculated from system variables
Examples of State Functions:
- Internal Energy (U)
- Enthalpy (H)
- Entropy (S)
- Gibbs Free Energy (G)
- Volume (V)
- Pressure (P)
- Temperature (T)
Properties of State Functions:
- Their values are independent of the path/process
- Only depend on the initial and final states of the system
- Any exact differential is a state function (dU, dH, dS, dG, etc.)
Not State Functions:
- Work (W) - depends on the path
- Heat (q) - depends on the path
- Other process quantities that depend on path taken
Significance:
- State functions are convenient thermodynamic bookkeeping tools
- Allow calculating energy changes between two states directly
- No need to consider the complicated path details
Key Point:
State functions have values determined entirely by the system’s current state, regardless of how that state was reached. This simplifies thermodynamic calculations.
Connect your understanding of potential energy to the new concept of bond dissociation enthalpy
Raising the PE to break a bond requires an input of KINETIC ENERGY—otherwise the bond will not break because it is at an energetically favorable position
U=internal energy of the system → changes in KE and PE add up to a change in U
Increases in KE of the system
Converts to PE, breaking bonds and/or deforming them away from minimum P.E/ lengths and angles
Some energy is lost as work, against external pressure, making the system expand (NOT IMPORTANT)
Heat entering the system (enthalpy) = internal energy + pressure x volume
H = U + pV
pV usually so small it is negligible
∆H > 0 for breaking bonds — energy added to the system — endothermic
∆H < 0 for forming bonds — energy leaves the system — exothermic
ALL OF THIS MUST BE TO A CONSTANT T [Wait for exothermic to cool down, or endothermic to get warm again]
Bond dissociation enthalpy is the heat input required to break one mole of bonds and maintain the same T at a constant pressure
Calculate the enthalpy of a reaction from bonds broken and formed
BOND DISSOCIATION ENTHALPY = total bonds broken (reactants) - total bonds formed (products)
Will determine if H is positive or negative
Enthalpy change does not depend on path—only initial and final states—IT IS A STATE FUNCTION
Regardless of forward or back it does not matter
Describe the difference between endothermic and exothermic reactions
Endothermic: takes energy in (immediately cold) to break bonds/create less strong bonds/configurations. KE converted to potential energy. (∆H > 0)
Exothermic: (immediately hot) releases potential energy as heat (kinetic energy) from the system when returned to a constant temperature—the energy of the products is less than that of the reactants (∆H < 0)
The new bonds formed/stronger bonds formed
Tend to be more likely spontaneous versus the other
double concentration -> collisions double -> successful reactions prolly double
collisions are independent , 2x collision with fixed fraction of energetic molecules means 2x reactions
T affects Reactivity through 1) fraction of molecules with sufficient energy 2) rate of collisions
Concentration affects reactivity through 1) rate of collisions
Temperature -
measure of Kinetic energy at molecular level
Heat-transfer of energy in the form of molecular (randomized ) KE
If heat energy is added to system where does it go? INCREASES Kinetic energy of system (raising temp)
change in KE/PE - change in internal energy “V”
CONVERTS to PE , breaking bonds and deformation of bonds
CONVERTS to work against external pressure -> system expands
Heat entering a system (at constant pressures) equals change in Enthalpy delta H
H = U + pV
enthalpy = internal energy + pressure/volume