Lectures 1-8 Flashcards
gases at room temperature
HEFONC & noble gases
1 L
1 dm^3
1 Pa
1 N/m^2
1 bar
1 x 10^5 Pa
1 kPa
1000 Pa
equal volumes of gas at constant T and p contain
equal number of particles
volumes of all gases extrapolate to zero at
0 K
R
8.314 J K^-1 mol^-1
V is proportional to
1/p
V is proportional to
T
V is proportional to
n
density
pM/RT
to find partial pressure
n(a) x (RT/V)
to find total pressure
(n(a) + n(b) + n(c)) x (RT/V)
to find mole fraction
n(a)/n(total)
to find mole fraction
p(a)/p(total)
to find partial pressure
x(a)/p(total)
kinetic molecular theory assumptions
negligible particle size, elastic collisions, no interaction
average kinetic energy of particles
3/2kT where k is R/N(A)
RMS speeds related to
temperature and molar mass
all gases at same temperature have same
kinetic energy, but not same average speed
effusion
gas escapes through a hole
diffusion
different gases mix
diffusion occurs from regions of
high to low concentration
rate of effusion is inversely proportional to
square root of molar mass
rate of effusion
number of molecules passing through a hole per second
particles attract each other
p decreases
particles repel each other
p increases
wavelength
distance from maximum to maximum
frequency
how many cycles a second
wavelength and frequency are
anti-proportional
wave number
inverse of wavelength; number of wavelengths per unit distance
light intensity is related to
amplitude
light colour is related to
wavelength
constructive interference
when two waves in phase interfere
destructive interference
when two waves out of phase interfere (amplitude = 0)
wavelength and energy content are
anti-proportional
same temperature
same kinetic energy
to calculate RMS, molar mass must be in
kg/mol
RMS
(3RT/M)^1/2
rate of effusion of gas 1/rate of effusion of gas 2
SQUARE ROOT OF: molar mass gas 2/molar mass gas 1
p1V1/T1 =
p2V2/T2
frequency =
c/wavelength
photons
energy packets of light
energy of a single photon (E)
hv
absorption spectroscopy
input light, analyse what was absorbed
in absorption spectroscopy, the sample transitions from a
lower to higher energy state
emission spectroscopy
input heat, analyse what was emitted
in emission spectroscopy, sample transitions from
higher to lower energy state
deltaE
E2-E1
predicting positions of lines in emission spectra
RH(1/n^2final - 1/n^2initial)
energy of emitted light
-hv or -hRydberg equation
quantum numbers
n(initial) and n(final)
emission quantum numbers
n(final) < n(initial) and deltaE is negative
absorption quantum numbers
n(final) > n(initial) and deltaE is positive
the Bohr model does not work because
emission spectra of many-electron atoms cannot be described, does not explain intensity of spectral lines, does not take into account particles as waves
photoelectric effect
electromagnetic radiation hits a metal surface and electrons are emitted
electrons are only ejected if light reaches
critical/threshold frequency
intensity increase causes
higher number of electrons to be ejected
photons must overcome both
binding energy and threshold frequency
kinetic energy of electron
hv(photon) - Ebind
kinetic energy of electron
hv(photon) - h x threshold frequency
Hz
1/s
de Broglie wavelength of matter
wavelength = h/p = h/mv
wavelength and mass are
anti-proportional
fast-moving particles have
shorter wavelengths
slow-moving particles have
longer wavelengths
Heisenberg’s Uncertainty Principle
it is not possible to accurately detect the exact position of a particle and its momentum at the same time
Heisenberg’s Uncertainty Principle
deltaxdeltap >/ h/4pi
wavefunction
used to describe the standing wave for an electron
for each energy state of an electron/particle, we assign a different
wavefunction
regions with positive and negative signs for each wavefunction
phases (there is a negative phase and a positive phase)
regions where wavefunction is zero
nodes
wavefunction^2
related to probability of finding a particle at a particular point in space
when wavefunctions are squared, nodes are
unchanged
point with maximum probability of finding a particle
maximum of wavefunction^2
point where probability of finding particle is zero
node
electron density
calculated through wavefunction^2, gives us probability of finding an electron in area around nucleus
orbital
a one-electron wavefunction in three-dimensional space
energy levels are always
negative
orbitals with the same energy are called
energetically degenerate
angular momentum quantum number; l
determines type & shape of orbital; tells us subshell/letter
limit for angular momentum quantum number (l)
n-1
magnetic momentum quantum number; ml
determines orientation of orbital; allows us to distinguish between each individual p, d orbital
limit to magnetic momentum quantum number (ml)
-l and l
principle quantum number; n
determines energy for that shell
s orbitals are
spherical
p orbitals are
dumbbell-shaped
two lobes of p orbitals have different
signs
two lobes of p orbitals separated by
angular node/nodal plane (zero)
keep radius fixed and
plot angles to explore shape
keep angles fixed and
play around with radius (can be graphed)
higher n,
higher number of nodes (for example, 1s orbital has no nodes, 2s orbital has one node, 3s orbital has two nodes)
the first time an orbital appears there are
no nodes
shell volume increases with
radius
wavefunction and probability decrease with
r
for larger shells, you’re more likely to find an electron
further away from the nucleus
to find wavelength when given mass
wavelength = h/mv
in many electron atoms, orbital degeneracy
disappears
electron spin is
quantised; two different angular momenta
spin magnetic quantum number ms for spin up/clockwise
+1/2
spin magnetic quantum number ms for spin down/anti-clockwise
-1/2
in an atom, no two electrons can have the same
four quantum numbers
the same orbital can only hold two electrons if they have
opposite spins
orbitals cannot hold more than
two electrons
fill up lower-lying orbitals
first
only fill next orbital once orbital below is
occupied
principal quantum number determines
row
superscripted number
number of electrons in that orbital
two degenerate orbitals want to be
singly occupied first
number of p orbitals in a shell
3 (6 electrons)
number of d orbitals in a shell
5 (10 electrons)
same group, same
electronic configuration in outer shell
core electrons
electrons in shells that are energetically lower than outer shell
nucleus has
positive charge Z
shielding is a
consequence of electron repulsion, which reduces net interaction between the positive nucleus and valence electrons
valence electrons do not experience full Z, only
effective charge Zeff
why do p orbitals have higher energies than s orbitals for many-electron atoms?
shielding
because 3d orbitals are shielded by 3s and 3p orbitals, they are
less stable and have higher energies than 4s orbitals
3d is filled after
4s
4d is filled after
5s
4f is filled after
6s
5d is filled after
4f
number of f orbitals in shell
7; 14 electrons
first transition metals
3d metals (fourth period)
4d transition metals appear after
5s orbital has been filled (fifth period)
chromium electronic configuration
[Ar]4s^13d^5
additional stabilisation in chromium is due to
half-filled 3d subshell with all 5 electrons having parallel spin
copper electronic configuration
[Ar]4s^13d^10
stabilisation in copper due to
fully-filled 3d subshell
after lanthanum, we fill
4f orbitals before remaining 5d orbitals
after actinium, we fill
5f orbitals before remaining 5d orbitals
after actinium, we fill
5f orbitals before remaining 5d orbitals
ionisation energy for a single H atom
13.6 eV
covalent radius
atomic radius estimated by halving distance between two chemically bound nuclei of the same type
size increases
down a group
as we go down a group,
principal quantum number (n) increases
size decreases
across a period
across a period, electrons from outer shell are drawn closer to nucleus due to
increasing positive charge
across a period, electron-electron repulsion
increases as we increase the number of electrons
cations are always
smaller than the parent atom
anions are always
larger than the parent atom
isoelectric
two species that contain the same number of electrons
for an isoelectric series, size decreases as
number of protons increases/nuclear charge increases
ionisation energy (IE)
energy required to remove the highest energy (most loosely held) electron from an atom/ion
removing an electron is
endothermic
first ionisation energy < second < third because
consecutively removed electrons cause a reduction in electron-electron repulsion and experience a greater attraction to Z
IE decreases
down a group
IE increases
across a period
across a period, nuclear charge
increases but the number of core electrons remains the same
high ionisation energy
small size
low ionisation energy
large size
electron affinity (EA)
the energy change associated with the addition of an electron to a gaseous atom
electron affinity is
exothermic, generally
the closer the vacant orbital is to the nucleus, the more favourable the addition of an electron and the
more negative EA; more energy released
EA decreases/is less negative
down a group
EA increases/is more negative
across a period
EA is not always exothermic due to
electron-electron repulsion
electronegativity
the ability of an atom to attract shared electrons to itself
electronegativity numerically
IE + EA
electronegativity increases
across a period
electronegativity decreases
down a group
metallic character increases
down a group
metallic character decreases
across a period
metals have relatively low
IEs
total number of orbitals in a shell
n^2
substance that has one or more unpaired electrons
paramagnetic
substance that has no unpaired electrons
diamagnetic
substance that has unpaired electrons that are aligned in a particular direction
ferromagnetic
