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Chemistry > Electronic Structure Of Atoms > Flashcards

Electronic Structure Of Atoms Flashcards

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1
Q

What is an energy level?

A

It’s a region of definite energy within an atom that electrons occupy

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2
Q

What elements have the same number of electrons in the outer most energy level

A

Are in the same group

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3
Q

How do elements emit light and give an example

A

When the atoms of the elements are supplied with energy under certain
conditions
Energy can be supplied by a flame

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4
Q

What flame colour does barium have?

A

Yellow-green

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5
Q

What flame colour does copper have?

A

Blue-green

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6
Q

What flame colour does lithium have?

A

Deep red

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7
Q

What flame colour does potassium have?

A

Lilac

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8
Q

What flame colour does sodium have?

A

Yellow

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9
Q

What flame colour does strontium have?

A

Red

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10
Q

Flame test

A
  • crush salt to be tested with mortar and pestle
  • dip soaked splint in salt and put splint in flame
  • note colour
  • repeat with other salts and note colour
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11
Q

Give example of flame tests in everyday life

A

Firework displays - strontium and barium

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12
Q

What is the discharge tube used for?

A

If an element is easily vaporised, it emits light of a characteristic colour when placed in discharge tube at low pressure and high voltage

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13
Q

What is a spectroscope?

A

It analyses light emitted by elements

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14
Q

What can flame tests be used for

A

To distinguish between different metals in their compounds

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15
Q

What is white light?

A

Mixture of visible light of all wavelengths

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16
Q

What occurs in a spectroscope when analysing white light?

A

The light waves are bent to different extents forming a continuous brand of colours or continuous spectrum [red to violet]

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17
Q

What is viewed when light emitted from a discharge tube containing hydrogen is analysed using a spectroscope?

A

A series of coloured lines of definite wavelength against a dark background is observed = line spectrum

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18
Q

How could you differ different elements?

A

By their emission spectrum. The emission spectrum of an element is characteristic of that element and is different from that of any other element

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19
Q

Why are certain colours not seen in the hydrogen/line spectrum?

A

Certain colours not able to be absorbed

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20
Q

What kind of light do elements emit?

A

Visible light and light in the ultraviolet and infrared regions

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21
Q

What does the hydrogen spectrum include?

A

Series of lines in both the visible and ultraviolet regions

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22
Q

What contains all the lines in the visible region of the hydrogen spectrum?

A

The Balmer series

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23
Q

What contains the lines in the ultraviolet region of the hydrogen spectrum?

A

Lyman series

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24
Q

What contains the lines in the infrared region of the hydrogen spectrum?

A

Paschen series

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25
Q

What does each element have?

A

A characteristic emission spectra and an absorption spectra

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26
Q

What is the absorption spectrum?

A

It’s the spectrum that is observe after the white light has been passed through an element

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27
Q

What does an absorption spectrum consist of?

A

A series of dark lines against a coloured background. The dark lines are at exactly the same wavelengths as the coloured lines in the emission spectrum of that element.

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28
Q

How can you measure the amount of an element in a sample

A

From the amount of light absorbed - atomic absorption spectrometer

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29
Q

The amount of light absorbed depends on what?

A

The amount of element in the sample

30
Q

How is the atomic absorption spectrometry useful?

A

Analysis of heavy metals - estimate the amount of lead in a blood sample

31
Q

What theory did Neil’s Bohr put forward to do with the line spectrum ?

A
  • Hydrogen electron is restricted to those regions of the atom that have certain energy values = energy levels and it cannot have an energy value lying between energy levels
  • when an electron moves from a higher energy level [E2] to a lower energy level [E1] a definite amount of energy is emitted, amount of energy emitted is equal to the energy difference = E2 - E1 = hf
  • atom absorbs amount of energy equal to the energy difference between e1 and e2 the electron will move to the higher level
  • each energy level has associated with it an integer called the principal quantum number
32
Q

What is the ground state?

A

It is the lowest energy state for the hydrogen atom

33
Q

What is the excited state?

A

It is a higher energy state for the hydrogen atom

34
Q

Describe the Bohr theory which accounts for the line spectrum of the hydrogen atom

A
  • Normal circumstances, the hydrogen electron is in n=1 energy level, the ground state -the lowest energy state- for the hydrogen atom.
  • if receives enough energy [eg, discharge tube], it moves to n=2, excited state -higher energy state- for the hydrogen atom.
  • it is unstable in this state and eventually will drop back to ground state, emitting energy equal to the difference between energies of n=2 and n=1 energy levels.
35
Q

How are the ultraviolet [lyman series] lines obtained?

A

When the electron falls from the n=3 to the n=1 level

36
Q

How do the visible lines in the hydrogen spectrum -the balmer lines- arise?

A

Electrons in n=3 energy level falls to n=2

37
Q

When electrons fall they emit energy or light at definite wavelength, what does the light appear as?

A

One of the line on the line spectrum of the element

38
Q

How is the line spectrum unique to every element?

A

In atoms of the element, the spacings between the energy levels are unique which give rise to unique electronic transitions

39
Q

What is the Heisenberg uncertainty principle?

A

It states that it is not possible to determine at the same time the exact position and velocity of an electron

40
Q

What is an atomic orbital?

A

It is a region in space where the probability of finding an electron is relatively high

41
Q

What are the three types of orbitals and what are their shapes

A

S orbitals -spherical shape
P orbitals - dumbbell shape
D orbitals - too complex

42
Q

How many electrons can all orbitals hold?

A

Two

43
Q

Of the orbitals, which have the highest and lowest energy?

A

3px, 3py, 3pz = equal energy
3s = lower energy
3d = highest energy

44
Q

In relation to orbitals, how does the energy impact the position?

A

The lower the energy, the closer it is to the nucleus

45
Q

What is an energy sublevel?

A

It is a subdivision of an energy level containing one or more orbitals, all of which have the same energy

46
Q

How many electrons can s sublevel hold?

A

2

47
Q

How many electrons can the p sublevel hold?

A

6

48
Q

How many electrons can the d sublevel hold?

A

10

49
Q

How many electrons can the f sublevel hold?

A

14

50
Q

What rules are used to weigh electrons to various sublevel and orbitals?

A

Aufbau principle - electrons will occupy the lowest energy sublevel available
Not more than two electrons can occupy an orbital at one time
Electrons occupy orbitals of equal energy singly where possible

51
Q

What are the limitations of the Bohr theory?

A
  • It only worked well for hydrogen
  • didn’t take into account that electrons have the properties of waves as well as particles
  • did not allow for Heisenberg uncertainty principle
  • did not explain discovery of sublevel
  • did not account for existence of orbitals
52
Q

What is the atomic radius of an element?

A

It is half the distance between the nuclei of two atoms of the element that are joined together by a single covalent bond

53
Q

What are the trends in the size of the atomic radius going down the group

A
  • electrons in inner levels partially neutralise attractive force of nucleus by repelling outer electrons
  • in group 1, an extra energy level is added and effected nuclear charge experienced by outermost electron is much less than full nuclear charge
  • atomic radius increases, due to the addition of extra energy levels, resulting in a extra screening by energy levels
54
Q

Trends in size of atomic radius going across a period

A
  1. Nuclear charge - exerts a greater attractive force on the outer electrons
  2. Screening effect - extra electrons are added to same energy level so there is no screening effect
55
Q

What is the first ionisation energy of an element?

A

It’s the minimum energy in kilojoules required to remove the most loosely bound electron from each isolated atom in a mole of the element in its ground state

56
Q

How is the first ionisation energy represented?

A

X = X+ + e-

57
Q

What is the first ionisation energy measured in

A

Kilojoules per mole

58
Q

Does the first ionisation energy increase or decrease across a period? And which is easier/harder?

A

The values increase across a period making it harder to remove the loosely bound electron

  • nuclear charge increases - strong force of attraction for the electrons from the nucleus and therefore electrons are held more tightly and so it’s harder to remove an electron
  • decrease in atomic radius due to the nuclear charge exerting an attractive force on outer electrons + no screening effect makes it easier to remove electron from an atom
59
Q

Does the ionisation energy increase or decrease going down a group? Is it harder or easier to remove loosely bound electron?

A

The energies decrease, easier

  • increase in atomic radius, easier to remove an electron despite increased nuclear charge
  • screening effect of inner energy levels, causing the nuclear charge experienced by the outermost electron being much less than the full nuclear charge
60
Q

Exception - beryllium and boron

A

Beryllium has higher ionisation energy value than boron. [full outer sublevel]
Extra electron is added to boron 2px orbital which has higher energy than 2s orbital. Electron is more readily removed than the 2s electron in beryllium, beryllium has a full outer sublevel and is stable
Nitrogen and oxygen

61
Q

What is the second ionisation energy?

A

It’s the minimum energy required to remove the most loosely bound electron from each singly charged positive ion in a mole of these ions

62
Q

How is the second ionisation energy represented

A

X+ = x2+ + e-

63
Q

Elements with similar outer electronic configurations but inner orbitals incomplete

A

Will have similar chemical properties

64
Q

Elements with different outer electronic configurations

A

Have different chemical properties

65
Q

What is the aufbau principle?

A

It states that electrons will occupy the lowest energy sublevel available

66
Q

What are the exceptions to electronic configurations?

A

Copper and chromium

[just take one from second last one and give to last]

67
Q

What are large and small atoms?

A

Large atomic radius - large atom

Small atomic radius - small atom

68
Q

Identify the main energy levels involved by in the electron transition that gives ride to Balmer series

A

3 and 2

69
Q

Why is it difficult to specify the absolute boundary of an atom?

A

Heisenberg uncertainty principle

Impossible measure position and velocity of electron in an atom

70
Q

Account for general trend in atomic radii going across the second period

A

Decreases

  • electrons in inner levels partially neutralise attractive force of nucleus by repelling outer electrons
  • the increasing nuclear charge exerts a greater attractive force on the outer electrons and pulls them in
  • electrons are added within the same energy levels, screening effect by inner levels is the same
71
Q

Set of successive ionisation energy values for electrons in carbon
[1086,2353,4620,6223,37831,47277]

How provide evidence for

  • no. electrons in a carbon atom
  • no. electrons in each main energy level in a carbon atom
A
  • each value = one electron

- gradual increase for first four values = four electrons in outer shell [easier to remove]

72
Q

Why is the first ionisation energy of oxygen lower than that of nitrogen despite the general increase

A

Nitrogen is relatively stable

-3 half filled 2p orbitals

Chemistry (15 decks)

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