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Physics Thermal
Physics Summer Y12
| Question | Answer |
|---|---|
| What is temperature? | A measure of the average kinetic energy of the particles in a substance |
| How does energy flow | From hot to cold always |
| How much K changes | 273 more |
| kelvin | No little circle |
| Useful phrase for this topic | Thermal equilibrium |
| Why thermometer into hot water inaccurate | Some E from water to thermometer |
| Brownian motion | Random motion of particle first shown with smoke particle in air (can also be shown with carbon particle in oil) |
| The conclusions we can draw from the motion of smoke particle in air | There are at any small time slightly more collisions in 1 direction than another Smoke p accel in this random direction This is evidence ps in gas/liquid move at random speeds in random directions Smoke ps much larger than air ps - only see smoke |
| Why we can only see smoke particles not air | Smoke particles about a micron, but air is about 10^-9, while light only ^-7 so generally cant see air as not large enough to block |
| What is the kinetic model? | Describes how all substances are made of atoms/molecules |
| How does the kinetic model describe solids? | Arranged in fixed regular lattice structure Vibrate around fixed point Have KE and a temperature |
| How does the kinetic model describe liquids? | Arranged close together randomly Move position around each other |
| How does the kinetic model describe gas? | Far apart, randomly distributed Can occupy the whole container Move with a range of speeds |
| What is internal energy | The sum of the *randomly distributed kinetic and potential energies of the *particles* in a substance IE = KE + PE KE tells us temp (T proportional to KE) PE tells us state of matter |
| Potential energy | Always negative technically. Solids have highest magnitude negative energy. PE of solids or liquids always less than zero, but ideal gases have zero potential energy (perfectly separated particles) Is due to particle separation |
| Why the flat sections on the heating/changing state graph | When flat can exist in either state (double point). If temp down, more change per second so more energy from PE to KE up to melting point. If temp up, less energy to KE as less transitioning per second, so lowers. |
| Specific heat capacity for states of matter | Not the same for each state |
| Gallium cooling | When temp reducing in a curve, goes through undercooling then sharply rises slightly before flattening as crystallisation happens (then curves down as it cools) |
| What is specific heat capacity? | Energy required to change temp of 1kg of a substance by 1K |
| Internal energy usage | ALWAYS MENTION. Also probably mention sum of KE and potential. |
| 0.8kg of copper at 76C in 2kg water at 15C. Calc final temp. | m1 c1 deltat1 = m2 c2 deltat2 m1 c1 (76 - final temp) = mc (final temp - 15) [+ m3c3etc] Arrange the final temp negatives like that to ensure both sides positive |
| Heat capacity | In JK^-1 Energy to heat whole sample, not per kg |
| Flat sections on the heating/changing state graph | Energy being put in is being transferred to potential energy |
| Specific latent heat | The energy needed to change the state of 1kg of a substance E = m L |
| SLH of fusion | SLH for a substance to change between solid and liquid phases |
| SLH of vaporisation | SLH for a substance to change between gaseous and liquid phases |
| How to use SLH in combo with specific heat capacity | Just add the energy needed to change state to the energy of heating |
| Specific latent heat practical | Kettle. Measure mass that boils (make sure to leave kettle until fully cool before measuring final. maybe you should only measure initial mass once already 100C?). Note down power rating of kettle and time heated for. |
| The stupid equation with 3 Ns | N = n * [N subscript A] Number of atoms/molecules = number of moles * Avogadro's number |
| temp relation to vibrations | Higher temp so higher amplitude of vibrations |
| What is an ideal gas? | A gas that obeys Boyle's Law |
| Theoretical assumptions for ideal gases | Large numbers of particles in constant, random motion No intermolecular forces All collisions are perfectly elastic Volume of molecules negligible to volume of gas |
| How gases act in practice | Simple molecules at everyday pressures and temperatures act like ideal gases e.g. O2 H2 N2, air, He at pressures up to 100s of atm Complex (e.g. CO2, CH4) deviate significantly at high temps and pressures (e.g. internal combustion engine) |
| How mass is used as a measure of a substance | Amount of substance in terms of inertia (for grav attraction, momentum, KE) |
| How moles are used as a measure of a substance | Amount of substance in terms of number of atoms/molecules (for rate of reaction, gas behaviour) |
| Mass of one mole | Equal to atomic mass in grams |
| Atomic mass | Remember to adjust depending on diatomic or monoatomic |
| Ideal gas equation | pV = nRT Pressure [Pa] x Volume = Number of moles x Universal Gas constant x Absolute temp |
| Universal gas constant value and unit | 8.31 J mol^-1 K^-1 |
| Gas laws all have what assumption | That the amount of moles of gas is constant |
| Boyle's Law (def need to remember name) | Ideal gas at constant temp (temp = boils/Boyles) P proportional to 1/V aka P1V1 = P2V2 |
| Charles' Law | Ideal gas at constant pressure V proportional to T aka V1/T1 = V2/T2 |
| Pressure Law aka Gay Lussac's Law | Ideal gas at constant volume P proportional to T aka P1/T1 = P2/T2 |
| General case for gas laws (most useful generally - works on most) | pV/T = cosntant so P1V1/T1 = P2V2/T2 |
| Atmospheric pressure | 1 x 10^5 Pa |
| Boyle's Law graphically | 1/x graph Depending on moles of gas the curve could be anywhere - these curved lines are called isotherms |
| Charles' and Gay Lussac's laws graphically | Directly proportional |
| GCSE explanation of pressure increasing with temp | Temp up so higher average KE so faster so more frequent and more average force per collision P = F/A and area constant |
| Difference between T and t | T generally for absolute temp on graph (K) t generally for non-absolute |
| Setting up for the ideal gas proof: What are the kinetic theory assumptions (slightly different to ideal gas) | Point molecules with negligible volume compared to the gas volume No intermolecular forces Molecules in continuous random motion All collisions perfectly elastic Time for collisions with container much shorter than time between collisions |
| Setting up for the ideal gas proof: Whats a root mean square? | A different type of average. It is the mean of the squares then squarerooted. It is always higher than the arithmetic mean. |
| Setting up for the ideal gas proof: Boltzmann constant | pV = nRT, where R is the universal gas constant k = Boltzmann constant = Universal gas constant/Avogadros number = R/(N subscript A), so R = [N subscript A]k |
| Setting up for the ideal gas proof: Alternate form of the ideal gas equation | k = R/(N subscript A), so R = [N subscript A]k So pV = n * NA * k * T moles x molecules per mole = number of molecules = N so pV = NkT Easier for certain questions depending on if Q says number of moles or molecules |
| Overview of ideal gas proof | Looking for pressure exerted by gas on box so need momentum for force |
| Ideal gas proof: working out change in momentum | Regular box with single molecule impacting wall with area A and travelling parallel to side with length d Momentum before collision with A = mu After collision = -mu, so change in momentum is -2mu (of wall on particle) |
| Ideal gas proof: working out force of particle on wall | Force = rate of change of p Time between collisions = 2d/u [taking avg pressure over time for scalability] Force of particle on wall = delta p/delta t = -(-2mu)/(2d/u) = 2mu^2/2d = mu^2/d It's -(-2mu) because particle on wall now so opposite |
| Ideal gas proof: working out pressure on wall due to 1 particle | Force due to 1 particle on face A = mu^2/d Pressure = F/A = (mu^2/d)/A = mu^2/V Therefore pV = mu^2 We now want to link this to pV = NkT |
| Ideal gas proof: working out pressure on walls when scaling up to a box of many randomly moving particles (at a variety of speeds) | Mechanically equivalent to 1/3 of the molecules moving in each plane at R.M.S. speed (like vector addition) For one molecule pV = mu^2 For N molecules, pV = 1/3 * N * m * (RMS)^2 = 1/3 * N * m * [c^2 with a bar above all of it = mean of squares] |
| Ideal gas proof: What is the mean KE of an ideal gas? | Mean KE = 0.5 x mass x mean of squares (again due to RMS being squared) |
| Ideal gas proof: linking mean KE of an ideal gas to the pressure exerted by the gas Let mean of squares just be c for this flashcard but always write it as C^2 with the bar above it all | Mean KE = 0.5 * m * c pV = 1/3 * N * m * c = NkT [1/3 * N * m * c = NkT] multiply both sides by 3/2 and cancel out N gives 1/2 * m * c = 3/2 * k * T |
| Ideal gas proof: What that equation is showing Let mean of squares just be c for this flashcard but always write it as C^2 with the bar above it all | We keep the 1/2mc instead of simplifying in order to show its the average KE of molecule (and also simplifies Qs where that is given) We've proven avg KE of ideal gas proportionality to temp |
| N2 relative atomic mass = 28 at 300K Finding out the average speed of the particles | [c^2 bar] isnt the answer, the RMS is |
| Symbol for RMS | C^2 bar inside a root symbol |
| Ideal gas follow-on proofs: Internal energy of an ideal gas | IE = KE + PE IE for one molecule = 0.5 x m x c^2 = 1.5 x k x T IE for one mole = Na x 1.5 x k x T = 1.5 x R x T IE for n moles of an ideal gas = 1.5 x n x R x T |
| Ideal gas follow-on proofs: Pressure and density relationship to speed | pV = 1/3 x N x m x c^2 p = 1/3 x (N x m x c^2)/V Nm = total mass, so Nm/V = density pressure = 1/3 x density x c^2 |
| What is the boltzmann distribution | Distribution of speeds in a sample of an ideal gas. speed on x axis and number of molecules on y. |
| Boltzmann distribution shape | Somewhat normal distribution shaped, but NOT one - shouldnt hit zero when x = 0 (some molecules will actually be stationary), and the tail end doesnt hit zero and stretches out Peak is modal speed, then slightly on is the mean speed, then RMS |
| Shape of Boltzmann as the temp increases | Spreads out as gas heats up (area isnt constant, just the sum of the points) as long as volume constant |
| Why are the pV = equations seen as energy equations | pV has units of joules |
| Defining specific heat capacity | Just use 'unit mass' and 'unit temperature rise' |
| Defining LHF and LHV | Say 'energy to change substance from X to Y' not just energy difference i think |
| Diagonal parts on a heating graph | PE still increasing, but KE also increasing |
| Supercooling curve | Supercools below freezing temp until nucleation begins. This crystalisation then releases bond energy, so it almost instantly rises to freezing point then stays flat for a while before curving down again. |
| Why cooling is always a curve | Newton's law of cooling is that it is based on difference with environment, so falls faster when hotter compared to ice bath |
| Lab method of investigating Boyle's Law | Syringe + weights Reducing pressure from atmospheric with more W. Use 1/V to find r as gradient will = 1/nRT |
| Boyle's Law Apparatus | Pump forces air into piston which applies pressure to coloured fluid (ignore the fluid in questions). This fluid compresses the air to a certain volume. Then there's just a reader for the pressure. |
| Why must absolute temp be used for the equations | Temp must be measured as an absolute meaning zero temp means zero KE of molecules hence zero pressure. This is kelvin |
| Why not much helium in the atmosphere compared to now | Always start with assumptions - assume behaves as ideal gas. Average KE and so RMS depend on absolute temp. Boltzmann distribution shows have range of speeds. So fastest atoms exceed escape v and escape. |
| Extra points for helium atmosphere hotter q | Earth hotter in past so rate of escape greater He nuclei alpha particles created by radioactive decay inside earth so low levels of He maintained - equilibrium levels between production and escape |
| Calorimeter | Calorimeter with known heat capacity (not specific) useful for liquid SHC experiment. has a built-in stirrer. |
Created by:
Pyrogearos2