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AS Chemistry > Chapter 3: Electronic Structure > Flashcards

Chapter 3: Electronic Structure Flashcards

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1
Q

What is electromagnetic spectrum

A

All the wavelengths of radiation in different forms of different energy energy

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2
Q

All waves travel at the same speed. How can they be distinguished

A

Wavelength

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3
Q

What is frequency

A

No of waves passing a point per second

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4
Q

Relate frequency and wavelength and energy

A

Shorter the wavelength the more the frequency and hence the more the energy

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5
Q

State wave equation

A

c=√f

velocity=wavelength×frequency

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6
Q

How do w, f and e vary with the spectrum

A

Decreasing from radio waves to gamma waves

Energy/freqeuncy increasing from radio to gamma

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7
Q

How is white light split into an atomic emmision spectrum

A

Wavelengths/frequencies are divided into colors by diffraction through a prism

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8
Q

What is line spectrum

A

White light split into different wavelengths/frequencies consisting of discrete lines

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9
Q

What happens when atoms are supplied with energy

A

Electrons gain energy move from lower(ground) to higher(excited state) energy levels

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10
Q

What happens as electrons gradually lose energy

A

Move from higher to lower energy levels and emit energy of particular wavelength

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11
Q

What prevents electrons from leaving the atom

A

Electrostatic force of attraction between oppositepy charged protons in nucleus and electrons in orbits

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12
Q

Why are neutrally charged neutrons important

A

Stability of elements with more than 1 proton, which would repel and nucleus falls apart

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13
Q

State the essentials of Bohr’s model

A

Electrons move in shells

For any atom there’s a fixed set of allowable orbits/stationary states

As long as electrons remain in their energy levels their energy is constant

An electron can only pass from one stationary state to another giving out/absorbing discrete/definite quantities of energy (quanta)

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14
Q

An electron moves up a level when energy is absorbed. Why does it eventually fall back and what is released

A

The excited electron is unstable
Electromagnetic radiation/packet of energy/photon

Photons of UV have more energy than of IR

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15
Q

Energy of photon emitted is =

A

Energy change in atom

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16
Q

Explain in depth the Planck equation (evidence of bohrs model)

A

Atoms emit photons of certain energies which give likes of certain frequencies as they can only occupy certain orbits

They can either be at this orbit or the other not in between meaning produce continuous energy changes.

Energy is changed only by discrete amounts aka quantized and line spectrum is quantized and not continuous

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17
Q

What happens when an electron is at it’s highest energy level

A

Atom is ionized

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18
Q

What is ionization energy

A

Energy needed to remove an electron from each atom in a mole of gaseous atoms

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19
Q

Why do lines with higher energies converge

A

Energy levels in the atoms are closer at higher energy

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20
Q

What foes quantum theory suggest

A

Electrons/particled having wave properties

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21
Q

What was wrong with Bohrs model

A

It applied well to hydrogen. Not other complex atoms. It assumes electrons trajectory can be precisely described. An electrons position cannot known by means of radiation as this will give it a kick and disturb the motion and send hurtling off in a random direction

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22
Q

What’s heiseinbergs uncertainty principle

A

We cannot know the position of an electron at a given moment of time
We can only predict where it’s likely to be ie spread in an electron cloud

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23
Q

What is an atomic orbital

A

Volume if space around nucleus in which there is maximum probability of finding a specified electron

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24
Q

The higher the energy of an electron the - it is from the nucleus. And - varies with energy

A

Further

Shape

25
Q

Energy levels have

A

Principle quantum numbers

26
Q

How many electrons in s sublevel

A

2

27
Q

How many electrons in p sublevel

A

6

28
Q

How many electrons in d sublevel

A

10

29
Q

How many electrons in f sublevel

A

14

30
Q

Expression for sub level maximum electrons?

A

2n^2

31
Q

State the order of orbitals

A

1s, 2s, 2p, 3s, 3p, 4s, 3p, 4p, 5s, 4d, 5p, 6s so on

32
Q

How is it that p sublevel onwards there are more than 2 electrons

A

There are 3 or more orbitals as in axis orientations

33
Q

What is paulis exclusion principle

A

Each orbital can hold 2 electrons so long as they have opposite spin which reduces effect of repulsion . This cases a small change in overall energy

34
Q

What is the Aufbau principle

A

ORBITALS ARE FILLED IN THE INCREASING ORDER OF THEIR ENERGIES

35
Q

What is hunds rule

A

When orbitals of identical energy are available electrons occupy these singly rather than in pairs as far as possible to reduce mutual repulsion

36
Q

Total energy of electron is determined by 4 factors:

A

Energy level
Sublevel
Slightly different energies of orbitals from Same principle levels due to magnetic field generated by moving electrons
Opposite spin

37
Q

Why are the outer most electrons lost first?

A

They are least attracted by the effective nuclear charge of protons

38
Q

What is the first ionization energy of an element

A

Energy needed to remove one electron from each atom in a mole of gaseous atoms to form one mole of gaseous singly charged ions (+)

39
Q

What is the second ionization energy of an element

A

Energy needed to remove one electron from each ion in a mole of gaseous singly charged ions to form one mole of gaseous double charged ions (2+)

40
Q

The successive ionization energy value _

A

Increases

41
Q

Why is it that successive ionization energy increases?

A

As each electron is removed the remaining ion becomes more positively charged. Fewer outer electrons are being attracted by the same no of protons. Successive electrons are closer to the nucleus and of lower energy. A strong force of attraction has to be overcome to remove these electrons

42
Q

What is the attraction between protons and electrons called

A

Electric field

43
Q

Electrons furthest from the nucleus are of

A

Highest energy

44
Q

Explain the 3 (+1) actors affecting ionization energy

A

•Size of nuclear charge
As no of protons increases as electrons are removed, +ve nuclear charge increases, and attractive force between nucleus and electrons increases so more (ionization) energy is needed to over come it.

•Atomic radius (distance from electrons to nucleus/half tge distance between nuclei of 2 closest atoms)
Size of atom depends on how strongly protons attract outermost electrons, and increases as shells increase.
Fα1/d^2 . Force of attraction between protons and electrons decreases as quantum shells/distance increase and the lower ionization energy

•Shielding effect
Electrons in filled inner shells repel those in outermost shells and reduce the effect of +ve nuclear charge. The net +ve charge outer electrons feel is called effective nuclear charge. The greater the shielding effect, the lower the ionization energy

(•Stability of electron configuration
More energy required to remove electrons from inert gases/full stable shell)

45
Q

Explain the ionization energy trend down a group

A

Decreases.
Down a group (period no increases) and the outer shell moves further away from nucleus and shielding effect increases due to increase in inner shells. So attraction for outer shell decreases and so does energy

46
Q

In the periodic table, group number denotes

A

Outer Shell electrons

47
Q

In the periodic table, period number denotes

A

Number of shells

48
Q

Explain the ionization energy trend across a period

A

Increases in general.
Effective nuclear charge increases as protons increase, but atomic radius stays the same as shell no remains constant. Attraction for outermost electron increases and energy also generally increases

49
Q

Name some elements which are an exception to the general trend across a period that energy increases

A

For boron, oxygen, aluminium, silicon, sulfur it decreases. Because of sublevels (group increases but sublevels are further than orbits)

50
Q

Give 2 exceptions to electron configuration that you know and why

A

Chromium 24
Should be: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d4 but is 3d5, 4s1
Copper 29
Should be: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d9 but is 3d10, 4s1
This is because eelectrons prefer to be single than paired so the available spaces in the 3d orbital are occupied by an electron from the 4s orbital to reduce repulsion and make it more stable while half filled

51
Q

The easier it is to remove an outer electron the _ it’s energy level

A

Higher

And vice versa

52
Q

Why is a logarithmic scale used

A

As electrons are lost the same protons hold other electrons more tightly hence increasing ionization energy

53
Q

In a graph, what 4 things can be defined?

A

Successive energies show the number of electrons in each level

Levels contain different numbers of electrons before they are full

Energies for each electron/level

By seeing the first large jump the no of valence electrons can be determined and hence the group (first electron is easier to remove and is thus in energy level)

54
Q

What are s, p d, f blocks?

A

Valence electrons of elements in these blocks are in these sublevels

55
Q

Which elements are in s, p and d blocks

A

S- groups 1 and 2 (eg ns2)
P- groups 3,4,5,6,7,0 (except He which is in S) (eg ns2np6)
D- transition (eg ns2np6nd8)

56
Q

How do graphs provide evidence for the existence of sublevels?

A

Graph is divided into subsections between noble gases, this corresponds to filling of one she’ll with electrons. Successive energies decrease even as groups increase (ie proton increases, same shell but energy is still less) this is due to order of sublevels that is, p is further than s in same energy level n

57
Q

Define atomic, neutron, nucleon/mass number

A

Atomic : no of protons in the nucleus
Neutron : no of neutrons
Nucleon/mass : no of protons and neutrons in a nucleus

58
Q

Write the 1st and 2nd ionization energy equations for an element X including state symbols

A

X(g)->X+(g) + e ΔHi1 ^(top) -O-
X+(g)->X++(g) + e ΔHi2 ^(top) -O-

-O- is standard conditions

AS Chemistry (11 decks)

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