Electronic Configuration And The Periodic Table

Question 0

What is Quantum Theory?

Answer 0

Matter can only emit do absorb energy in small fixed amounts

Question 1

What is an atomic spectra caused by?

Answer 1

Electrons moving between diff energy levels - fixed for any 1 atom. We say that the energy of electrons in atoms is quantised

Question 2

What happens when an electron in an atom absorbs a photon of energy?

Answer 2

It moves from a lower energy level to a higher one. When the electron drops back down, energy is emitted.

Question 3

When are atoms said to be ‘excited’ ?

Answer 3

When they absorb energy and emission spectra arise from the movement of electrons from a higher to lower energy when the excited atom returns to its ‘ground state’

Question 4

What does the frequency of the line in the emission spectrum represent?

Answer 4

The difference in energy between the levels. We call these energy levels shells/sub shells.

Question 5

What does each line in an emission spectrum represent?

Answer 5

Radiation of a specific wavelength/frequency from which these differences in energy can be calculated

Question 6

What is the ground state?

Answer 6

The lowest possible electronic configuration the electrons in an atom can adopt

Question 7

What represents the ionisation energy of an electron?

Answer 7

An electron escaping from level n=1 to infinity which corresponds to the electron breaking away from the atom completely

Question 8

What happens to the energy levels as energy increases?

Answer 8

They come closer together until they converge.

Question 9

What does the difference in energy between the ground state and convergence limit correspond to?

Answer 9

The energy required for the electron to break away from the atom

Question 10

How do you find the energy of a photon?

Answer 10

The difference between the 2 energy levels ( deltaE= h(nu))

Question 11

What is the principal Quantum number?

Answer 11

n.
Determines the main energy level.
It can have values n=1,2,3etc (n=1 is nearest the nucleus)
The numbers deter me the size and energy of the shell. As n increases the potential energy and distance from nucleus increases

Question 12

What is the second Quantum number?

Answer 12

l .

Determines the shape of the sub shell and is labelled as s,p,f,d. This can have values from 0 to (n-1)

Question 13

What is Heisenberg’s uncertainty principle?

Answer 13

It is impossible to state precisely the position and the momentum of an electron at the same instant. So it isn’t possible to define a point in space where the electron is certain to be found and it is therefore necessary to define regions in space where the probability of finding an electron is high.

Question 14

What are atomic orbitals?

Answer 14

The volume in space where the probability of finding an electron is greater than 90%. The overall size of each orbital is governed by n, actual shape governed by l

Question 15

What is the third Quantum number?

Answer 15

m.

Relates to the orientation in space of the orbital. Dependant on l as m can take any whole value between -l and +l

Question 16

How many electrons can each atomic orbital hold?

Answer 16

2

Question 17

Describe s orbitals

Answer 17

Spherical with size and energy increasing as n increases.

Question 18

Describe p orbitals.

Answer 18

Have a value of l=1 and therefore there are 3 possible orientations in space, corresponding to m=-1,0,+1. The 3 p orbitals are degenerate (have the same energy as each other) and the same shape, dumbbell shaped at right angles to one another. Each orbital defined as if it lies along a set of x,y,z axes. The 2p orbitals are thus 2px,2py,2pz. The 3p orbitals would be the same but larger and at higher energy.

Question 19

Describe d orbitals.

Answer 19

Occur in five different orientations corresponding to the third quantum number m=+2,+1,0,-1,-2. Important in the examination of properties of transition metals.

Question 20

What is the fourth quantum number?

Answer 20

The spin. Has values of +0.5/-0.5. In any orbital containing two electrons they must be paired, with spins opposed.

Question 21

What is the Aufbau principle?

Answer 21

When electrons are placed into orbitals the energy levels are filled up in order of increasing energy, e.g. 1s before 2s.

Question 22

What is the Pauli exclusion principle?

Answer 22

An orbital cannot contain more than 2 electrons and they must have opposite spins. E.g. [⬆️⬆️] isn’t allowed! must be [⬆️⬇️]

Question 23

What is Hund’s rule?

Answer 23

When there are degenerate orbitals in a subshell electrons fill each one singly with spins parallel before pairing occurs. E.g. [⬆️ ][⬆️ ][⬆️ ]

Question 24

What are the relative energy corresponding to each orbital in increasing order?

Answer 24

1s,2s,2p,3s,3p,4s,3d,4p,5s etc.

Question 25

What are 5 factors which affect ionisation energy?

Answer 25

1) atomic size- greater atomic radius=further outermost electron is from nucleus and ionisation energy decreases.
2) nuclear charge- more protons=more difficult to remove an electron
3) shielding effect
4) subshells. E.g. Boron IE is lower than beryllium because boron involves taking one electron from 2p, but beryllium involves taking 1 electron from the full 2s subshell. Since full subshells are relatively stable, IE is greater.
5) half full subshells require more energy as they’re relatively stable.

Question 26

What is electronic configuration?

Answer 26

The arrangement of electrons in energy levels and orbitals of an atom

Question 27

What causes the spectral lines in an emission spectrum?

Answer 27

When an excited electron falls from a higher energy level to its ground state, releasing a photon of energy. The difference in energy relates to a specific frequency/wavelength which produces lines of that frequency/wavelength to appear as an emission spectrum