Atomic Structure

Question 1

Chemical reactivity is determined by:

Answer 1

The movement of electrons between atoms.

Question 2

Thomson (1898)

Answer 2

Plum pudding model: mainly positive with areas of negative electrons embedded within.

Question 3

Rutherford (1909)

Answer 3

Electrons are arranged equidistant from the nucleus with spherical orbits

Question 4

Rutherford weaknesses

Answer 4

Suggests all electrons are equal, makes it difficult to explain atomic emission spectra.

Question 5

Electron excitation

Answer 5

Heat/light energy is supplied to an atom, causing them to move from a ground to an excited state. This state is less stable, so electrons don’t remain there long, and when they return to lower energy states, they emit electromagnetic radiation - typically a photon of visible or UV light.

Question 6

Photon Energy

Answer 6

The energy of the emitted photon is calculated using the difference between the energy of the two levels.

Question 7

Atomic Emission Spectrum

Answer 7

The emitted light is dispersed into a spectrum using a prism or diffraction grating. Spectrum is discrete and unique to each element.

Question 8

Energy Gaps

Answer 8

Fixed and unique to each element - the larger the energy gap between two levels, the higher the energy of the emitted photon.

Question 9

Example: Hydrogen emission spectra

Answer 9

When hydrogen atoms are excited, the emitted light can be divided into specific series: Lyman (transitions to n=1 are in the UV region), Balmer (n=2 in the visible light region), Paschen (n=3 infrared region).

Question 10

Photon energy equation

Answer 10

E = hv = hc/λ

Question 11

Photon energy relevance to AE

Answer 11

shows the link between the energy of a photon and its frequency.

Question 12

Bohr (1913)

Answer 12

Similarly structured to Rutherford, however, electrons are at different distances to the nucleus and exist in quantised (discrete) orbits. Electrons closer to the nucleus are more tightly bound and therefore are more stable and have lower energy. E- can absorb energy and be excited or relaxed to different energy levels; this produces AE spectra.

Question 13

Bohr’s atom weaknesses

Answer 13

• Atomic spectra measured in a magnetic field are different to those measured outside
• Works for hydrogen but less so with heavy atoms
• Doesn’t explain the periodic properties of the periodic table

Question 14

De Broglie’s relationship

Answer 14

Relates pure particle to pure wave properties using Einstein’s E = mc^2 (particles) and Planck’s E = hc/λ (waves); finding that if those two statements are correct, then h/λ = mc = p (momentum)

Question 15

Heisenberg uncertainty principle

Answer 15

States that though with particles we can determine position and momentum exactly, this isn’t possible for wave/particles.

Question 16

Observed electron behaviour

Answer 16

When observing an electron in the classical sense, lightwaves reflect off the surface of the electron and reach the eye. These light waves add energy to the electron and change its momentum; therefore, observed electrons inherently behave differently.

Question 17

Implications of wave particle duality for atoms

Answer 17

An electron in orbit must have a definitive position and an exactly defined path; therefore, if electrons do not have an exact position or path, they cannot exist in orbits.

Question 18

Orbitals

Answer 18

A region of space within which an electron resides, although the exact location is unknown. The orbital represents a stationary wavefunction: the probability of finding an electron within a certain region.

Question 19

Wave mechanics of electron orbits

Answer 19

Electron orbit circumference should be an integral multiple of its wavelength to ensure that the wave meets in phase after completing one full circle. If it were out of phase, the wave would cancel itself out.

Question 20

Schrodinger’s equation

Answer 20

Describes how the quantum state of a physical system changes over time—how the wavefunction changes and therefore where electrons could be found. Shows how energy from movement and attraction impacts the wavefunction.

Question 21

Wavefunction Ψ

Answer 21

A component of Schrodinger’s equation that gives insight into the possible places or behaviours of a particle.

Question 22

Probability density Ψ^2

Answer 22

Ψ^2(x,y,z) is the probability of finding an electron at (x,y,z).

Question 23

Implications of Schrodinger’s equation

Answer 23

Describes electron behaviour, predicts location, helps to determine where energy levels are.

Question 24

Principle quantum number (n)

Answer 24

Determines the size of an orbital, distance from the nucleus, and therefore the energy

Question 25

Angular momentum number (l)

Answer 25

Determines the shape of an orbital, all orbitals with the same l value will have the same shape. Only the hydrogen atom energy levels are unaffected by the value of l.

Question 26

Why are H energy levels unaffected by l?

Answer 26

The attractive forces between the proton and electron are spherically symmetrical; this means that regardless of the value of l, all orbitals with the same n value experience the same potential from the nucleus. H atoms also do not experience electron-electron interactions; therefore energy is determined solely by distance from the nucleus.

Question 27

Magnetic quantum number (ml)

Answer 27

Determines the orientation of the orbital rlative to Cartesian axes.

Question 28

Naming orbitals

Answer 28

As the quantum numbers describe solutions to Schrodinger's equation for hydrogen, they also provide probability distributions and energy states. Orbitals are therefore named after the quantum numbers that describe them.

Question 29

Radial wavefunction (R)

Answer 29

The function of the distance of an electron from the nucleus. The radial aspect refers to how the probability changes as you move further away from the nucleus. relates to n and l.

Question 30

Radial distribution functions

Answer 30

shows how likely an electron is to be found at specific distances from the nucleus. The maximum point shows the most probable distance. RDF = probability of electron x number of locations at said distance.

Question 31

Radial nodes

Answer 31

points at which probability is 0; they appear in higher energy levels. For n, the number of radial nodes is n-l-1.

Question 32

Consequences of RDF

Answer 32

- The area of maximum probability increases with increasing n but is independent of l. - There is a small probability of finding the electron a very long way from the nucleus; hard to accurately define orbital size. - Nodes exist in further out energy levels

Question 33

Angular wavefunction (Y)

Answer 33

Dependent only on the l and ml numbers, it describes how the probability varies with angular direction around the shape of the orbital. Determines the shape and orientation of an orbital.

Question 34

Link between l and ml

Answer 34

For every value of l, ml can take any integer between -l and l. e.g. when l = 1 there are 3 possible values of ml and therefore 3 possible orbital shapes.

Question 35

Weaknesses of Shrodinger's equation

Answer 35

Potential energy (V) includes nucleus-electron attraction and electron-electron repulsion, which is 0 for H atoms. For atoms with >1 electron, the probability distribution must be determined before the potential energy - making the equation impossible to solve

Question 36

Mitigating SE weakness

Answer 36

H atom orbital wavefunctions can be used for many-electron atoms. Orbitals ar the same shape as the hydrogen atom. HOWEVER, o energies are dependent on both n and l.

Question 37

Orbital energies many-electrons

Answer 37

Electrons in orbitals with a larger n have higher energy as they have a probability distribution further away from the nucleus and are therefore less tightly bound.

Question 38

Shielding effect

Answer 38

Electrons closer to the nucleus prevent the outer electrons from feeling the full nuclear charge that they would if they were isolated. These electrons therefore have a lower effective nuclear charge and higher energy than expected.

Question 39

Why are orbital energies s < p < d < f?

Answer 39

When superimposing 2s and 2p RDFs, it is seen that the residual maximum of 2s is closer to the nucleus; therefore, 2s electrons on the whole have greater nuclear charge and lower energy levels.

Question 40

How Z affects orbital energies (3d, 4s)

Answer 40

As Z increases, the more penetrating 4s orbitals will experience greater nuclear charge and lower energy. 3d orbitals are closer to the nucleus and will also have lower energy. In K and Ca, 4s orbitals are more penetrating than 3d orbitals and have lower energy as a result. From Sc onwards, vice versa.

Question 41

Filling orbitals

Answer 41

Electrons fill orbitals starting from those with the lowest energies. With some exceptions, the 4s subshell tends to fill before the 3d.

Question 42

Aufbau principle

Answer 42

In the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy level before occupying higher-energy levels

Question 43

Spin quantum number (ms)

Answer 43

Describes the intrinsic spin (angular momentum) of an electron and has two possible values: 1/2 (spin up, or -1/2 (spin down)

Question 44

Pauli exclusion principle

Answer 44

No two electrons in any given system can have identical values for all four quantum numbers.

Question 45

Hund's rule

Answer 45

Electrons will occupy degenerate orbitals (of the same energy) singly and with parallel spins before pairing up. Makes for the most stable state.