Unit 2 Test: Gases and Thermochemistry
ideal gas conditions
high temperature, low pressure, and low number of moles (low number of moles = molecules are less crowded = more space to move freely)
calorimetry: hot metal in water
metal will lose thermal energy until H2O equilibrium is reached (hot metal is causing temperature change) for H₂O: q=energy absorbed by water (+), m=mass of H₂O, c=4.18Jg⁻¹°C⁻¹, ∆T=temperature change of H₂O for metal: q=energy released by metal (-), m=mass of metal, c=specific heat of metal, ∆T=temperature change of metal
total pressure of a gas mixture
sum of pressures (partial pressures) of individual gases that make up gas mixture
Maxwell-Boltzmann distribution
the distribution of energies (and therefore speeds) of the molecules in a gas
Avogadro's law
- as the number of moles (n) increases, the volume most increase b/c increasing the number of particles increases the number of collisions and in order to maintain constant pressure, the volume must increase - V1/n1=V2/n2 - only holds true at low pressure and high temperature (ideal condition)
Boyle's Law
- if the volume is decreased, the gas particles will hit the wall more often, thus increasing the pressure (temperature constant) - the graph of P versus V is hyperbolic - for a given quantity of a gas at constant temperature, the product of pressure and volume is a constant - P1V1=P2V2 - ideal gases are expected to have constant values of PV - law breaks down at high pressures /low temperatures/non-ideal conditions (molecule no longer free to move)
non-ideal gas conditions
- low temperature: gas molecules slow down and become more likely to condense - high pressure: gas molecules are more likely to stick together - even though we assume gas particles have no volume, they do; the volume available is actually less than 100% of the container size - IMFs cause pressure to be lower than predicted with ideal gas law (attraction between molecules causes fewer collisions and therefore lower pressure)
vapor pressure
- pressure created by molecules that strike the walls of a container when escaping into the gaseous phase - each type of liquid has its own specific vapor pressure - doesn't vary with the surface area of the liquid or size of the sample - varies with temperature; as temperature rises vapor pressure increases (as temperature rises, KE energy increases meaning more molecules are transitioning to the vapor stage, more molecules creates higher pressure) - whenever H₂O(l) is present, H₂O(g) is present also
overall enthalpy change for a reaction
- sum of the bond breaking and bond making - there is an endothermic and an exothermic component to every reaction - ∆Hrxn: energy exchange per one mol of substance reacted, often expressed in kJ/mol
Charles' Law
- when a gas is heated the particles speed up (particles hit the walls more often and with more force) so they only way to keep the pressure constant is to increase the volume of the container - V1/T1=V2/T2 - the graph of temperature versus volume is linear - gases with strong intermolecular forces between particles require more heat to expand because more energy is needed to break the bonds
using bond energy to calculate ∆Hrxn
1. draw Lewis dot diagram 2. count the bonds of each type & apply the bond values w/ appropriate signs 3. sum the net energy change 4. answer represents the energy change that would occur per mole of rxn (rxn is 'run one time')
molar volume at STP
22.4 L
ideal gas law
PV = nRT (R=(L x atm)/(mol x K))
speed of particles in a sample of gas
RANGE of speeds
Kinetic Molecular Theory (KMT)
a model used to explain gas behavior w/ the ASSUMPTION that gases act ideally 1. gas particles do not take up any space in a container (volume of 0) 2. gas particles are always in motion, the pressure of a gas changes when (1) how often the gas particles strike the walls of the container changes or when (2) how forcefully they strike the walls changes 3. particles have no attractive or repulsive forces between them; collisions are perfectly elastic 4. the only thing that changes the Kinetic energy of gas molecules is a change in temperature 5. the pressure exerted by a gas does NOT depend on the identity of the gas; N2, O2, H2, etc.... all cause the same amount of pressure in a vessel assuming that temperature and volume of the container remain the same
particle speeds at higher temperatures
at higher temperatures the average speed of the particles is higher because the MAJORITY of the particles have sped up, however NOT ALL of them have → there is more distribution of speeds present among the sample → peak of distribution curve shifts to the right and gets shallower because there is a greater range of speed (peak has same area as other curves b/c the number of molecules is not changing)
How would gas behavior change is the number of moles of the gas was doubled but everything else is constant?
average speed: no change b/c temperature and molar mass are constant KE: no change b/c temperature is constant pressure: higher w/ more moles b/c more collissions mean free path: lower with more moles b/c less distance between particles
using heating and cooling curves to calculate ∆Hrxn
calculate ∆Hrxn for each phase and phase change 1. solid: q=mC∆T for specific heat and temperature change of solid phase 2. melting/fusion: q=∆Hfus x amount of substance 3. liquid: q=mC∆T for specific heat and temperature change of liquid phase 4. vaporization/evaporation: q=∆Hvap x amount of substance 5. gas: q=mC∆T for specific heat and temperature change of gaseous phase
calculating the molar mass from density
density: g/v 1. moles=g/MM 2. PV=nRT → PV=(g/MM)RT 3. MM=(gRT)/(PV) → MM=(densityRT)/P..."molar mass kitty cat": good cats put dirt over their p
bond energy
energy required to break one mole of a certain type of bond (given as positive values in textbook tables)
potential chemical energy
energy stores in bonds between atoms AND attractive forces between molecules
intermolecular forces
forces of attraction between molecules (gas laws and states depend on IMFs)
Hess's Law
from a list of reactions and their ∆Hrxn's, multiply and/or flip each reaction so that each compound not needed in the final reaction is eliminated (present on both sides) and add the ∆Hrxn's to find the total ∆Hrxn for the final reaction
Are product bonds stronger in an exothermic or endothermic reaction?
in an exothermic reaction, the product bonds will be collectively stronger than the reactant bonds because no additional energy is needed to be absorbed for them to form (opposite is true for endothermic reaction)
Why does raising the temperature of a reactant mixture often speed up the reaction?
increasing the temperature increases the average speed and KE of the particles, meaning that more particles are now moving the speed needed for the reaction to occur (more particles reaching activation energy)
endothermic reaction (+)
more energy is absorbed from surroundings than what is released (products lie higher than reactants)
exothermic reaction (-)
more energy is released than absorbed (products lie lower than the reactant)
relationship between mole fraction and total/partial pressure
partial pressure/total pressure = mol/total mol Xi x Pt = Pi
thermal energy transfer
q=mC∆T
forming bonds ALWAYS
releases energy (-)/exothermic
breaking bonds ALWAYS
requires energy (+)/endothermic
coffee cup calorimeter
rxn in an aqueous solution in which the solute changes the solution's temperature as it dissolves q=energy transfer, m=mass of solute+solvent (entire solution), c=specific heat of solution, ∆T=temperature change of solution
∆H°f values
show energy exchange caused by a special type of reaction, formation rxns (forming a single species from common reactants), often given in a textbook table
What is the only factor that changes kinetic energy of a gas?
temperature; as it increases the average KE (and particle speed) increases b/c each individual molecule is given more energy
heating water by combusting something underneath it
the combustion of the alcohol causes the temperature to of water to change q=energy transfer, m=mass of water, c=4.18.18Jg⁻¹°C⁻¹, ∆T=temperature change of water
particle speed based on molar mass
the higher the molar mass, the slower the average particle speed/velocity (at constant temperature) b/c heavier objects move slower
mole fraction
the number of moles of a particular substance expressed as a fraction of the total number of moles (Xi)
What happens to PV when T is held constant?
the product of PV must be constant
calculating ∆Hrxn using ∆H°f values from a table
∆Hrxn=∑∆H°fproduct - ∑∆H°freactants
