Week 1 Lecture 1 Basic Concepts Flashcards
Name & define three type of systems
- Open - everything can leave
- Isolated - nothing can leave
- Closed - matter can’t leave but energy can (inc. thermal energy)
Name & define three types of properties of a system
- Extensive - depend on the size of the system
- Intensive - independent of the size of the system
- Specific - extensive properties divided by mass (lowercase)
What is an independent property?
One that can be varies while the other property is held constant
Define phase (in thermodynamics)
A region of space where matter is static in it’s properties - e.g done mixing ect.
Define thermodynamic equilibrium
All three types of equilibrium present:
- Thermal equilibrium - no temperature difference between system & it’s surroundings
- Mechanical equilibrium - forces are balanced
- Material equilibrium - means phase & reaction eq (i.e materials are done mixing & reacting or ‘transport’ & ‘ converting species’)
Define process
A change that a system undergoes from one equilibriums state to another (via process paths)
Define quasi-equilibrium
Literally an ‘almost’ equilibrium: when a process is slow enough the conditions of the system can be assumed to be (approx.) constant
What is the Celsius scale based off of?
The boiling point and freezing point of water
How is absolute zero found?
- By saying some quantity is proportional to temperature
- Finding the constant in terms of a known point in Kelvin (triple point of water)
- Plotting values on a graph & extrapolating back
Can anything be equal to absolute zero?
No; it can only tend towards it (because nothing is ever ideal)
What is the oth law of thermodynamics?
It’s about implied thermodynamic equilibrium; “Went two objects are separately in thermodynamic equilibrium and with a third object, they are in equilibrium with each other”
Name three pressure equations
- P=f/a
- P=\roh gh and integral form from modelling layers of fluid
- The aeroplane pressure equation (pressure at an altitude exponential equation) - p=p_0e^{-GPE/KbT}
Name & define two types of pressure
- Absolute pressure - measured relative to vacuum (as is anything ‘absolute’)
- Gauge pressure - measured relative to atmosphere (what you would see on a pressure gauge)
Define energy
Ability to do work (where total energy is sum of micro and macroscopic energy)
What is the difference between microscopic and macroscopic energy?
Microscopic energy is w.r.t the outside reference frame, macroscopic energy is w.r.t the molecules inside the system
Define internal energy (of an ideal gas)
The sum of the KE and PE of gas molecules (where PE=U=U_trans+U_vib+U_rot+U_p)
Why is potential energy for a molecule positive at (relatively) short distances?
Because repulsive forces do work to move the molecules apart
Name ideal gas laws
Boyle - P, V
Charles - V, T
Avogadro - V, n
Gay Lussac - P, T
When is the ideal gas law not valid?
At low temperatures and huge pressures
What is a PVT plot?
A 3D plot of P vs V vs T (with an exponentially decreasing shape)
Define partial pressure
It’s the pressure exerted by one gas in a mix of gasses if you took all of the other gases out (so in the same volume)
Define Dalton’s Law of additive pressures & Amagat’s Law of additive volumes
For gas mixes you add the properties:
Dalton: the pressure of the gas mix is equal to the sum of the pressures of the individual gases
Amagat: the volume of the gas mix is equal to the sum of the volumes of the individual gases
What is a mole fraction?
- For gas mixes the no. of moles of one gas / no. of moles of the total gas mix
- It’s the same as the number pressure fraction and volume fraction (as is unitless)
State archimedes principle
The buoyant force is equal to the weight of the fluid displaced
