2. Atomic Structure Flashcards

2.1 The nuclear atom 2.2 Electron configuraiton 2.3 Ionisation Energy

1
Q

Define isoelectronic

A

Atoms/ions that have the same number of electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Define isotonic

A

Atoms/ions that have the same number of neutrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Define isotopes

A

Atoms of the same element with a different number of neutrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

What are the characteristics of isotopes?

A

Same chemical properties but different physical properties

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Reading a mass spectrum: Define the x & y axes

A

x: Mass/charge ratio
y: Relative abundance

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Relationship between Energy & Wavelength

A

As E increases, wavelength decreases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Relationship between Energy & Frequency

A

As E increases, frequency increases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

What is the purpose of the emission spectrum?

A

It is evidence for the existence of electrons in discrete energy levels which converge at higher energies

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Electromagnetic region associated with n=1

A

Ultraviolet

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Electromagnetic region associated with n=2

A

Visible light

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Electromagnetic region associated with n=3

A

Infrared

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

What is the convergence limit?

A

The gap between the lines decrease until each series converges to a limit

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Describe the Bohr model in terms of energy levels, sub-levels, orbitals

A

Each energy level (1, 2, 3, …) has different numbers of sub-levels. Each sub-level has different numbers of orbitals, depending on which sub-level. Each orbital holds only 2 electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Shapes of orbitals

A

s: spherical
p: dumb-bell

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

What is ground state?

A

The lowest energy state

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What are the 3 rules of electron arrangement in orbitals?

A
  1. Aufbau principle
  2. Hund’s rule
  3. Pauli’s exclusion principle
17
Q

Define Aufbau principle

A

Electrons are added progressively to the orbitals starting with lowest energy
(Remember the arrow diagram)

18
Q

Define Pauli Exclusion Principle

A

Paired electrons can only be stable when they spin in opposite directions so that the magnetic attraction which results from their opposite spins counterbalances the electrical repulsion

19
Q

Define Hund’s Rule

A

When filling up a sub-level, each orbital must be occupied singly before they are occupied in pairs

20
Q

Exceptions to the Aufbau principle

A

Chromium & Copper

21
Q

How are electrons lost when forming cations

A

Lost from the orbital with the highest energy

22
Q

How are electrons added when forming anions

A

Added to the vacant orbital of highest energy

23
Q

Define 1st ionisation energy

A

The miniumum energy required in removing one mole of valence electrons from one mole of gaseous atoms to form 1 mole of singly positively charged gaseous ion

24
Q

What are the 2 factors affecting ionisation energy?

A
  1. Nuclear charge

2. Shielding Effect

25
Explain in detail how nuclear charge affects ionisation energy
The greater the nuclear charge, the greater the electrostatic force of attraction between the positively charged nucleus and negatively charged electrons, therefore the greater the ionisation energy
26
Explain in detail how the shielding effect by inner electrons work
Electrons in inner shells repel valence electrons, increasing the shielding effect of the electrostatic forces of attraction of the nucleus on the valence electron, therefore lowering ionisation energy
27
Explain why successive ionisation energies of an atom increase with the removal of each electron
An increasing amount of energy is required to remove successive electrons from an increasingly positive ion due to increasing electrostatic forces of attraction between nucleus and the valence electrons
28
General trends in 1st ionisation energy
Increases across period, decreases down a group