Gases and Intermolecular Forces Flashcards
overview of gases
- atoms and molecules are generally not in contact
- intermolecular bonding can be observed
pressure
- gas molecules exert pressure on their container through collisions
- bigger than 0 pascals, the molecules are in motion
- collective result of collisions is pressure (p) which is a macroscopic property
equation for pressure
p = F/A
force in newtons (N)
unit area in square meters (m2)
atmospheric pressure
weight of a column of air molecules in the atmosphere above an object
Boyle’s law and equations
if you decrease volume, you increase pressure
if you increase volume, you decrease pressure
PV = k (constant)
P1V1 = P2V2
what do you need to do to correct for pressure gauges?
pressure gauges measure above atmospheric pressure so add 1 atm to internal pressure of a tire etc.
Charles’ law
as gas heats up, it expands
V is proportional to T
V1/T1 = V2/T2
avagadro’s law
as the number of moles of gas increases, the volume increases (constant P and T).
limitation: only works with ideal gas laws (can’t keep things constant)
intramolecular force
attractions and interactions between atoms within a molecule
intermolecular forces
forces between molecules (relatively weak)
dispersion forces
fluctuating, temporary dipole that results in relatively weak electrostatic attraction. applies to all atoms and molecules
instantaneous dipole
electrons (in motion) are distributed asymmetrically for an instant
induced dipole
instantaneous dipole in one atom distorts the electrons in another. they form dispersion forces
trends in dispersion forces
- larger/heavier atoms and molecules have stronger dispersion forces (I2 is solid but F2 is a gas)
- as atoms/molecules increase in size and mass, melting and boiling points increase (more heat needed)
- forces strengthen as surface area increases (long)
polarizability
how easily an electrostatic charge can distort a molecule’s charge distribution
dipole-dipole attraction
electrostatic force between the partially positive end of one polar molecule and the partially negative end of another
- high boiling point than non polar molecules
- can be liquids at room temp
hydrogen bonding
intermediate forces between hydrogen of one molecule and a more electronegative atom on another (FON)
- group 15, 16 and 17 hydrides increase boiling points down the groups (dipole-dipole become weaker instead)
- lightest of the hydrides should have lower boiling point but don’t (hydrogen bonding!)
intermolecular forces of halogens
- form diatomic molecules such as H2
- single covalent bonds
- end on overlap of p orbitals
compare Cl2, F2, B2 and I2
Cl2 and F2 expand and contract in response to pressure, move freely, collide - intermolecular forces are small (gases)
Br2 move freely but not as freely (less room), cannot be compressed, doesn’t expand a lot, intermolecular forces are stronger so is liquid
I2 has more electrons than Br2, little space between molecules, doesn’t expand or contract much, ordered in arrays and forces are strong enough to prevent molecules from sliding past each other. can vibrate (solid).
reasons:
- large KE molecules remain seperate as a gas
- intermolecular forces are large and force molecules close as a liquid or gas
ideal gas law with Pa or kPa
PV = nRT
P = Pa V = m3 n = mol R = 8.314J/mol/k T = k 1000L in m3 so can use kPa and L instead
ideal gas law with atm
PV = nRT
P = atm V = L n = mol R = 0.08206L/atm/mol/k
rearrangement of ideal gas law
M = mRT/PV
density
varies with conditions
p (ro) = m/V
p = m/V = PM/RT
rules of density
- density increases as pressure increases
- density decreases with increasing temperature
- density increases with M
why can CO2 extinguish fire?
- higher molecular weight than air - denser
- displaces oxygen around the fire
what does the ideal gas law assume?
- gas has negligible forces between its constituent atoms and molecules
- negligible volumes
this is not true for real gases!!
van der Waals equation for real gases
(p + n^2a/V)(V-nb) = nRT
corrects volume and negligible forces
- gases can be liquified!!
why is PV/nRT less than 1 at low pressure?
intermolecular attractions - lower pressure than expected
why is PV/nRT greater than 1 at high pressure?
gases do have a volume - finite volume of gas atoms/molecules
