Energy Flashcards
1
Q
Thermodynamics
A
- The study of energy transfer that we can measure in the macroscopic words
- Does not make any assumptions about the microscopic world (like KMT)
- Heat transfers until thermal equilibrium is established
- ΔT measures energy transfer
2
Q
Energy
A
- The capacity to do work or produce heat
- Most energy can be exchanged between objects through contact
3
Q
Work
A
Form of energy based on moving an object over a distance
4
Q
Heat
A
- The form of energy that flows betweem 2 objects because of their difference in temperature
- Quantity of hotness (mass + speed)
- Thermal energy
5
Q
Temperature
A
Average kinetic energy
6
Q
What is a calorie
A
- heat required to raise ttemp of 1.00 g of H2O by 1.0 °C
7
Q
Units of energy
- cal –> kcal
- kcal –> Cal
- cal –> joules
A
- 1000 cal = 1 kcal
- 1 kcal = 1 Calorie (a food “calorie”)
- 1 cal = 4.184 joules
8
Q
First law of thermodynamics
A
The total energy content of the univerws is constant
9
Q
Potential energy
A
Stored energy that an object has by virtue of its position or compsotion
10
Q
Kinetic energy
A
energy of motion or energy of a reaction
11
Q
5 forms of energy (PE or KE)
A
- Electrical - KE
- Heat / thermal energy - KE
- Light / radiant energy - KE
- Nuclear - PE
- Chemical - PE
12
Q
Internal energy
A
- sum of all forms of KE (motion of atoms/molecules) and PE (rep by tthe chemical bonds and IMFs) in a system
- E = internal energy
- E = PE + KE
13
Q
System
A
The object under study
14
Q
Surroundings
A
Everything outside the system
15
Q
Direcionality of heat transfer
A
- Heat always transfers from hoter objec to cooler one
16
Q
Exothermic reaction
A
- heat transfers from system to surroundings
- PE of products < PE of reactants
17
Q
Endothermic reaction
A
- heat transfers from surroundings to system
- PE of products > PE of reactants
18
Q
Change in energy equation
A
- ΔE = q + w
- change in energy = heat + work
- if isolated system (closed) then ΔE of system = 0
19
Q
q, w, ΔE signs (+ or -)
A
- +q: system gains energy
- -q: system releases heat
- +w: system gains energy from work (compression)
- -w: system releases energy by doing work (expansion)
- +ΔE: system gains energy
- -ΔE: system releases energy
20
Q
work equation
A
- w = -PΔV
- constant pressure
21
Q
Heat capacity (c)
A
- The heat required to raise an object’s temp by 1 C
- Cup vs tub: tub has higher heatt capacity
22
Q
Specific heat
A
- The amount of heat required to raise the temp of one gram of a substance by one degree celsius
- H2O: 4.184 J / gºC
- How much energy is tranferred due to temp difference?
- The heat (q) lost or gained is related to
- sample mass
- change in T
- specific heat capacity
23
Q
specific heat equation: q=
A
q = smΔT
24
Q
Calorimetry
A
- The amount of heat absorbed or released during a physical or chemical change can be measured, usually by the change in temp of a known quantity of water in a calorimeter
25
Enthalpy
* A state function that accounts for heat flow in the processes (chemical or physical) that occurs at constant pressure when **no forms of work** are performed except P-V work
* No gas, no change in volume
* H = ΔE + PΔV
* ΔH = q (at constant P)
* For a chemical reaction:
* ΔH = Hproducts - Hreactants
26
State function
* A property that is independent of the pathway
* Enthalpy
* Temperature
* Energy
27
Not state functions
* Heat and work
28
Rate
change in property / time
29
3 things must happen to molecules in order for a reaction to occur
* Collide
* Energy to react (have to collide with enough force)
* Orientation
30
Activated complex
Top of hill on energy diagram
31
Rate determining step
Slowest reaction is the RDS
32
Exothermic fast vs slow reaction
* Fast: small upward hill
* Slow: large upward hill
