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Flashcards in Electronic Structure Of Atoms Deck (72):
1

What is an energy level?

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

2

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

Are in the same group

3

How do elements emit light and give an example

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

4

What flame colour does barium have?

Yellow-green

5

What flame colour does copper have?

Blue-green

6

What flame colour does lithium have?

Deep red

7

What flame colour does potassium have?

Lilac

8

What flame colour does sodium have?

Yellow

9

What flame colour does strontium have?

Red

10

Flame test

-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

11

Give example of flame tests in everyday life

Firework displays - strontium and barium

12

What is the discharge tube used for?

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

13

What is a spectroscope?

It analyses light emitted by elements

14

What can flame tests be used for

To distinguish between different metals in their compounds

15

What is white light?

Mixture of visible light of all wavelengths

16

What occurs in a spectroscope when analysing white light?

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

17

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

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

18

How could you differ different elements?

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

19

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

Certain colours not able to be absorbed

20

What kind of light do elements emit?

Visible light and light in the ultraviolet and infrared regions

21

What does the hydrogen spectrum include?

Series of lines in both the visible and ultraviolet regions

22

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

The Balmer series

23

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

Lyman series

24

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

Paschen series

25

What does each element have?

A characteristic emission spectra and an absorption spectra

26

What is the absorption spectrum?

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

27

What does an absorption spectrum consist of?

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.

28

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

From the amount of light absorbed - atomic absorption spectrometer

29

The amount of light absorbed depends on what?

The amount of element in the sample

30

How is the atomic absorption spectrometry useful?

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

31

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

-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

What is the ground state?

It is the lowest energy state for the hydrogen atom

33

What is the excited state?

It is a higher energy state for the hydrogen atom

34

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

-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

How are the ultraviolet [lyman series] lines obtained?

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

36

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

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

37

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

One of the line on the line spectrum of the element

38

How is the line spectrum unique to every element?

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

39

What is the Heisenberg uncertainty principle?

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

40

What is an atomic orbital?

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

41

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

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

42

How many electrons can all orbitals hold?

Two

43

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

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

44

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

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

45

What is an energy sublevel?

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

46

How many electrons can s sublevel hold?

2

47

How many electrons can the p sublevel hold?

6

48

How many electrons can the d sublevel hold?

10

49

How many electrons can the f sublevel hold?

14

50

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

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

What are the limitations of the Bohr theory?

-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

What is the atomic radius of an element?

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

53

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

-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

Trends in size of atomic radius going across a period

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

What is the first ionisation energy of an element?

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

How is the first ionisation energy represented?

X = X+ + e-

57

What is the first ionisation energy measured in

Kilojoules per mole

58

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

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

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

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

Exception - beryllium and boron

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

What is the second ionisation energy?

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

How is the second ionisation energy represented

X+ = x2+ + e-

63

Elements with similar outer electronic configurations but inner orbitals incomplete

Will have similar chemical properties

64

Elements with different outer electronic configurations

Have different chemical properties

65

What is the aufbau principle?

It states that electrons will occupy the lowest energy sublevel available

66

What are the exceptions to electronic configurations?

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

67

What are large and small atoms?

Large atomic radius - large atom
Small atomic radius - small atom

68

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

3 and 2

69

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

Heisenberg uncertainty principle
Impossible measure position and velocity of electron in an atom

70

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

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

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

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

72

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

Nitrogen is relatively stable
-3 half filled 2p orbitals