Energy

Question 1

What is energy, and the equation for energy transfer?

Answer 1

• Energy is the capacity to do work or supply heat (1st law)
• ΔE = q + w
  
  • A change in energy = heat and work combined
• You get ΔE if you measure heat exchange at constant volume. (bomb calorimetry)
  
  • ΔE = q_V

Question 2

What are the two types of energy?

Answer 2

• Potential energy - the energy of position or properties
  
  • Chemical energy is a form of this, because of the electrical attractions and repulsions within atoms, molecules and ions
• Kinetic energy - energy of motion
  
  • Can be work (ex: moving pistol) or heat (ex: making molecules move faster)

Question 3

What is internal energy?

Answer 3

• The sum of molecular kinetic and potential energies
• Heat and work
• Extensive property - higher with more molecules

Question 4

Temperature vs heat

And the equation that relates them

Answer 4

• Temperature is an intensive property, determines the direction of heat flow
• Heat is a process, the transfer of energy from an object with high temperature to one at low temperature. Extensive property.
  
  • Conducted by molecular collisions

q = m • c • Δ t

heat flow = mass • heat capcity or specific heat • temp change

Question 5

Bond energy

breaking vs. formation

Answer 5

• Breaking bonds requires energy, consumes it.
  
  • Higher final potential energy.
• Forming bonds generates/releases energy
  
  • Lower final potential energy, more stable. The rest left as kinetic energy, which can be work or heat.

Question 6

Maxwell-Boltzmann distribution of speeds

Answer 6

• If you raise the temperature of your system, the speed of the molecules in that system increases
• If all of the molecules in your system have a larger speed and are colliding with each other, they are more likely to react with each other

Question 7

What is reaction energy?

Answer 7

Reaction energy is the sum of all the bonds being broken and all the bonds being formed.

Question 8

Exothermic reactions

Answer 8

• The system gives off heat to the surroundings
  
  • Release more energy during the reaction than has been consumed
• Overall bond energy is more stable at the end (since bond formation releases energy)
• Ex: combustion
• Breaking bonds is exothermic
• We use this term when discussing enthalpy.
  
  • Negative enthalpy of reaction means an exothermic reaction
  • (For energy, we use the terms exergonic and endergonic.)

Question 9

Endothermic reactions

Answer 9

• The system absorbs heat from surroundings
• It takes more energy to break the bonds than you get at the end, so it pulls energy from its surroundings
• Breaking bonds is endothermic
• We use this term when discussing enthalpy.
  • (For energy, we use the terms exergonic and endergonic.)
• ΔH of your reaction (your system) is positive, then the reaction was endothermic

Question 10

What is our main equation for heat transfer,

and what do the variables stand for?

Answer 10

q = msΔT

can also be written

q = nCΔT

q = heat (energy units, Joules)

m = mass (g)

n = moles

s = specific heat capacity (when working with grams) J/g° C

C = heat capacity (when working with moles) J/moles° C

Question 11

What does the sign of q tell us?

Answer 11

• If q is a positive value→ the system gains energy. Heat moves from the SURROUNDINGS into the SYSTEM.
  
  • Endothermic.
• If q is a negative value→ the system loses energy. Heat is a product.
  
  • Exothermic

Question 12

What is work, and the main context in which we use it?

Answer 12

• Work is movement against a restraining force
• Context: expansion of a gas. Look for creation of gas molecules
  
  • The gas is moving against pressure, but increasing in volume
  • w = -PΔV
  • P pressure is always positive

ΔV is always positive for expansion of a gas, because growing volume

If “w” work is negative, that means your system has done work

* depends on Δn<sub>gas</sub>

Question 13

What is enthalpy?

Answer 13

• Enthalpy = energy without the work. Represented by H.
  
  • ΔH = ΔE - w
• Usually a good stand in for energy. Same unit, Joules (J).
• You get ΔH when you measure the heat exchanged at constant pressure. (Solution calorimetry)
  
  • ΔH = q_P
• If the ΔH of your reaction (your system) is positive, then your reaction is endothermic.
  
  • Don’t forget to first find q of your system: In an isolated system, q_system = -q_surroundings

Question 14

What are the two kinds of calorimetry?

Answer 14

• Bomb calorimetry
  
  • Constrain the volume to force the system to transfer all energy as heat rather than work
  • ΔE = q_V (at a constant volume)
• Solution calorimetry
  
  • In a cup. Atmospheric pressure is the constant. Not measuring work, so define Enthalpy as energy without the work.
    
    • ΔH = q_P (at a constant pressure)

Question 15

State function vs path function

Answer 15

• State functions just depend on final # - initial #.
  
  • Ex: energy and enthalpy
  • We do use signs because that tells us the direction of the flow.
• Path functions depend on how you got there
  
  • Ex: heat and work. Transfers, not amounts.

Question 16

What is Hess’s law, and when do we use it?

Answer 16

• Allows us to manipulate state functions
  
  • The total enthalpy change for the reaction is the sum of all changes
  • Can do this with enthalpy or energy
  • Multiply by coefficients, by -1 to reverse, add up the enthalpies at the end