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A Level Chemistry-CIE > Atomic Structure > Flashcards

Atomic Structure Flashcards

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1
Q

What is the relative mass and relative charge of a proton?

A

relative mass: 1
relative charge: 1+
(These are not the actual charges and masses. They are charges and masses compared with each other in a simple ratio)

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2
Q

What is the relative mass and relative charge of a neutron?

A

Relative mass: 1
Relative charge: 0
(These are not the actual charges and masses. They are charges and masses compared with each other in a simple ratio)

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3
Q

What is the relative mass and relative charge of an electron?

A

Relative mass: 1/1836
Relative charge: 1-
(These are not the actual charges and masses. They are charges and masses compared with each other in a simple ratio)

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4
Q

Describe the distribution of mass within an atom.

A

The mass is concentrated at the nucleus of an atom.

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5
Q

What is the nucleus made up of?

A

The nucleus is made up of particles called nucleons. There are two types of nucleons: protons and neutrons.

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6
Q

Describe the behaviour of beams of protons, neutrons and electrons in an electrical field.

A

When subatomic particles are passed between two oppositely charged plates the protons will be deflected on a curved path toward the negative plate because, they are positive. Whereas the electrons will be deflected on a curved path towards the positive plate because they are negative. Neutrons will continue on a straight path because they have no charge.

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7
Q

What is the atomic number(z)?

A

The atomic number is equal to the total number of protons present in the nucleus of an atom. Atoms of the same element have the same atomic number and number of protons.

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8
Q

What is the mass number (A)?

A

The mass number/nucleon number (A) is equal to the total number of neutrons and protons in an atom.

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9
Q

How do you find the neutrons in an atom given the nuclied notation?

A

The number of neutrons in an atom= mass number-atomic number. = a - z.

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10
Q

How do you find the neutrons in an atom given the nuclide notation?

A

The number of neutrons in an atom= mass number-atomic number. = a - z.

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11
Q

What is an ion?

A

An ion is a charged atom i.e. the number of protons is not equal to the number of electrons.

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12
Q

When is a positive ion formed?

A

Positive ions are formed by the loss of electrons. The +ive on ions is equal to the number of electrons lost.

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13
Q

When is a negative ion formed?

A

Negative ions are formed by the gain of electrons. The -ve charge is equal to the number of electrons gained.

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14
Q

What is an isotope?

A

Isotopes are atoms with the same atomic number and a different mass number. Isotopes are atoms with the same number of protons but a different number of neutrons.

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15
Q

Do isotopes have similar or different chemical properties and why?

A

Isotopes have similar chemical properties because they have the same number of electrons.

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16
Q

Do isotopes have similar or different physical properties and why?

A

Isotopes have different physical properties such as small differences in density because they have different masses/ different number of neutrons.

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17
Q

Which sub-atomic particle is responsible for the chemical properties of the atom?

A

The electrons are responsible for the chemical properties.

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18
Q

What are two ways isotopes can be represented?

A

isotopes can be represented by nuclide notation. Chemists also name them by emitting the proton number and placing the nucleon number after the name for example the isotopes of hydrogen can be called:
hydrogen-1
hydrogen-2
hydrogen-3

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19
Q

What are the three naturally occurring isotopes of hydrogen?

A

The three naturally occurring isotopes of hydrogen are:
1. Protium (atomic number 1 and mass number 1)
2. Deuterium (one proton, one neutron, and one electron)
3. Tritium (one proton and two neutrons)

They all have the same number of protons.
Protium or common hydrogen is the only atom without a neutron in its nucleus.

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20
Q

What is the definition of cathode?

A

The negatively charged electrode in electrolysis is called the cathode. Positively charged ions move towards the cathode.

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21
Q

What is the definition of anode?

A

The positively charged electrode is called the anode. Negatively charged ions move towards the anode.

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22
Q

What is the definition of a nucleon number?

A

The nucleon number (A) is the number of protons plus neutrons in the nucleus of an atom. This is also known as the mass number.

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23
Q

What is a principal quantum number (n)?

A

A principal quantum number (n) represents the shell that the electrons occupy. These principal quantum numbers are numbered according to how far they are from the nucleus. The lowest energy level, n=1, is closest to the nucleus, the energy level n=2 is further out, the third shell is n=3 and the 4th shell is n=4.

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24
Q

What is the relationship between principal quantum numbers and energy?

A

The larger the principal quantum number, the higher the energy and the further the shell is from the nucleus.

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25
Q

What is the definition of electron configuration?

A

The arrangement of electrons in an atom are called its electron configuration.

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26
Q

What does the electron configuration show?

A

Electron configuration shows the number of electrons and types of orbitals in each energy level.

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27
Q

What is the formula used to calculate the maximum number of electrons each principal quantum shell can hold?

A

2n^2.
The first shell n=1 can hold a maximum of 2 electrons.
The second shell n=2 can hold a maximum of 8 electrons.
The third shell n=3 can hold a maximum of 18 electrons.
The fourth shell n=4 can hold a maximum of 32 electrons.

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28
Q

What is an atomic orbital?

A

Each subshell contains one or more atomic orbitals. An atomic orbital is a three-dimensional region of space around the nucleus of an atom that can be occupied by one or two electrons where the probability of finding the electrons is maximum.

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29
Q

What does electronic structure refer to?

A

Electronic structure refers to how electrons are present around the nucleus in different shells (energy levels), sub-shells/ sub energy levels and orbitals.

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30
Q

What are sub-shells/ sub energy levels?

A

A subshell is specific type of orbitals in a shell (e.g. the p subshell contains 3 p orbitals). the shells/ energy levels further consist of sub-shells (sub energy levels). These sub shells/ sub energy levels are s, p, d and f.

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30
Q

What is the order of energy of the electrons in the subshell in any principal quantum shell?

A

In any principal quantum shell, the energy of electrons the energy of electrons in the subshells increases in the order s<p<d<f.

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31
Q

How can you figure out how many sub shells are there in a particular atom?

A

The number of sub-shells in a shell= shell number (N).

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32
Q

How many orbitals does the s sub-shell contain?

A

S subshell has only one orbital. (One orientation)

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33
Q

How many orbitals does the p sub-shell contain?

A

P subshell contains 3 orbitals. (Three orientations).

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34
Q

How many orbitals does the d sub-shell contain?

A

D subshell has five orbitals. (Five orientations).

35
Q

How many orbitals does the f sub-shell contain?

A

F subshell has seven orbitals. (Seven orientations).

36
Q

How many electrons can a single orbital accommodate?

A

A single orbital, no matter what type can accommodate a maximum of 2 electrons.

37
Q

What are the maximum number of electrons represented as?

A

The maximum number of electrons are represented by the power of the sub-shell. For example, S subshell can accommodate a maximum of 2 electrons and therefore s has the power of 2 (s^2).

38
Q

Describe the s-orbital?

A

An s-orbital has a spherical shape (like a soccer ball). The s-orbital has no specific direction. i.e. it is non-directional. There is one s orbital in each shell from n=1 upwards (a total of two s electrons per shell.

39
Q

Describe the p-orbital?

A

A p-orbital has an hourglass shape or a dumb-bell shape with two lobes. There are three p orbitals in each shell from n=2 upwards ( a total of six p electrons per shell). The p orbital has higher energy than s.

40
Q

What does the electron density of the p-orbital consist of?

A

Its electron density consists of two lobes.

41
Q

What representation can be used to show electrons in subshells along with energy levels?

A

1s^1
where the first 1 represents the principal quantum number.
The letter represents the subshell.
The power 1 represents the number of electrons.

42
Q

Describe the d orbital.

A

Five d orbitals in each shell from n=3 upwards (a total of 10 d electrons per shell), higher energy than p.

43
Q

What are the three p orbitals called?

A

The three p orbitals are called degenerate orbitals (meaning same energy).

44
Q

What is the order in which the atomic orbitals are filled?

A

The order in which atomic orbitals are filled depends on their relative energies. The lowest available energy level is filled first.

45
Q

What are the three principles on which electron configuration is based on?

A
  1. Aufbau Principle
    2.Pauli’s exclusion principle
  2. Hund’s rule
46
Q

Explain the Aufbau principle:

A

This principle states that electrons always fill the lowest energy orbitals first.

47
Q

Explain the Pauli’s exclusion principle.

A

This principle states that an orbital can have a maximum of two electrons with opposite spins.

48
Q

Why do orbitals have opposite spins when they are paired in an orbital?

A

The electrons are negatively charged so they repel each other. By spinning in opposite directions, a magnetic attraction develops helping the electrons gain stability and reducing the repulsion created by their negative charges.

49
Q

What is the Hunds rule principle?

A

If a set of degenerate orbitals (having same energy) is available, then the electrons will prefer to stay apart with unpaired spin in separate orbitals.

50
Q

What is the stability of atomic orbitals- principle?

A

This rule is only applicable to chromium with 24 electrons and copper 29 electrons. This rule states that completely filled atomic orbital is more stable with less energy than half filled or partially filled orbitals.
For example:
P6= Completely filled atomic orbital (less energy) MOST STABLE.
P3: Half-filled atomic orbital.
P1, P2, P4, P5: incomplete or partially filled atomic orbital. LEAST STABLE.

51
Q

What are the two ways to right down electronic configuration?

A
  1. s, p, d, f configuration.
  2. Using spin-in-box diagram.
52
Q

What is the electronic configuration order for transition elements?

A

1s→2s→2p→3s→3p→4s→3d→4p→5s→4d→5p

53
Q

Why does the electron configuration of chromium and copper not follow the usual pattern?

A
54
Q

How should you write the electronic configuration for positive ions?

A

When a positive ion is formed electrons are lost from the outermost shell or valence shell.
For example: magnesium with 12 electrons will have the electron configuration of 1s^2 2s^2 2p^6 3s^2. There are two electrons in its valence shell.
magnesium ion (+2) has the electron configuration of 1s^2 2s^2 2p^6 as the two electrons from the valence shell have been lost.
Positive or negative ions always have a completely filled valence shell.

55
Q

How should you write the electron configuration for negative ions?

A

When a negative ion is formed electrons are gained.
For example: oxygen with 8 electrons has the electronic configuration 1s^2 2s^2 2p^4 (1s^2 2s^2 2px^2 2py^1 2pz^1).
Oxygen ion (-2) has the electron configuration 1s^2 2s^2 2p^6.
Positive or negative ions always have a completely filled valence shell.

56
Q

How should we write the electron configuration for transition metal ions?

A

Electrons are first lost from the 4s orbitals and then the 3d orbitals.

57
Q

What is the relationship between orbitals and the periodic table?

A

The arrangement of elements in the periodic table reflects the electronic structure of elements. The periodic table is split into “blocks of elements”.

58
Q

Which elements are the s block elements and why?

A

The elements in group 1 and 2 have their valence (outermost) electrons in an s subshell. These are therefore called the s-block elements.

59
Q

Which elements are called the p block elements and why?

A

The elements in group 13-18 (except helium) have their valence (outermost) electrons in a p subshell. These are therefore called the p block elements.

60
Q

Which elements are called the d block elements and why?

A

The elements that add electrons to the d subshells are called the d block elements. Most of these are transition elements.

61
Q

What is the definition of a free radical?

A

A free radical is defined as a molecular species which can contain an unpaired electron in its atomic orbital and can exist independently.

62
Q

How are free radicals formed?

A

Free radicals are formed from the absorption of energy.

63
Q

Why are free radicals highly reactive?

A

Free radicals are highly reactive as they have high energy.

64
Q

What are the free radical electrons represented as?

A

The unpaired electron is shown as a dot on the symbol for example: chlorines free radical is shown as Cl.
Group of atoms can also be free radicals for example CH3 has a carbon atom with an unpaired electron.

65
Q

What is the definition of the first ionisation energy?

A

The first ionisation energy of an element is the energy needed to remove one electron from each atom in one mole of atoms of the element in gaseous state to form one mole of gaseous 1+ ions.

66
Q

What is the symbol for the first ionisation energy?

A

I.E1 (1 standing for the first) or ΔHi.

67
Q

What are the units for ionisation energy?

A

The units are kJ mol-1

68
Q

What is the definition of second ionisation energy?

A

The second ionisation energy is the energy required when one mole of unipositive gaseous ions are converted into one mole of dipositive gaseous ions (+2 charged ions) by the loss of one electron.

69
Q

Is ionisation energy an endothermic or exothermic process?

A

Ionisation energy is positive for neutral atoms meaning that ionisation energy is an endothermic process. This is because energy must be supplied to overcome the electrostatic attractive force between the nucleus and the electron.

70
Q

Why is the second ionisation energy always greater than the first ionisation energy?

A

For each element the successive ionisation energies increase. This is because the charge of the ion gets greater as each electron is removed. As each electron is removed there is a greater attractive force between the positively charged protons in the nucleus and the remaining negatively charged electrons. Therefore more energy is needed to overcome these attractive forces.

71
Q

What are the four factors which affect ionisation energy?

A
  1. The size of the nuclear charge
  2. Distance of outer electrons from the nucleus.
  3. Shielding effect of inner electrons.
  4. Spin-pair repulsion.
72
Q

What is the trend between ionisation energy and nuclear charge?

A

As the nuclear charge increases so too does ionisation energy increase.

73
Q

What is the trend between ionisation energy and nuclear charge?

A

As the nuclear charge increases so too does ionisation energy increase.

74
Q

Explain the trend between ionisation energy and nuclear charge.

A

As the number of protons increases the positive nuclear charge increases. This increase in nuclear charge exerts a greater electrostatic attraction on the electrons and therefore more energy is required to remove electrons.

75
Q

What is the trend between ionisation energy and distance of outer electrons from the nucleus?

A

The further the outer electron shell is from the nucleus the lower the ionisation energy.

76
Q

Explain the trend between ionisation energy and the distance of outer electrons from the nucleus.

A

The force of attraction between positive and negative charges decreases rapidly as the distance between them increase. So electrons in shells further away from the nucleus are less attracted to the nucleus than those closer to the nucleus.

77
Q

What is trend between ionisation energy and the shielding effect of the inner electrons?

A

The greater the shielding effect of outer electrons by the inner electron shells the lower the ionisation energy.

78
Q

Explain the trend in ionisation energy and the shielding effect of inner electrons.

A

The greater the shielding of outer electrons by the inner electron shells the lower the attractive forces between the nucleus and the electrons hence less energy is required to remove electrons.

79
Q

Explain the relationship between ionisation energy and spin pair repulsion.

A

Electron pairs in the same atomic orbital in a subshell repel each other more than electrons in different atomic orbitals. This increased repulsion results in less energy needed to remove an electron.

80
Q

Explain the trend in ionisation energy across a period.

A

Across a period from left to right, the ionisation energy increases as with each successive element one proton is added so the nucleus is becoming more positive. This increase in nuclear charge exerts a greater electrostatic attraction on the electrons and therefore more energy is required to remove electrons.

81
Q

Explain the trend in ionisation energy down a period.

A

Going down a group, the ionisation energy decreases. This is because with each successive element there is an extra occupied energy level so the electron being removed is further away from the nucleus. Also, this electron is more shielded from the positive charge of the nucleus, by the extra inner occupied energy levels. This increased distance and shielding of the outer electrons from the nucleus, makes the electrostatic attraction weaker and electrons are more easily removed.

82
Q

Why is the ionisation energy of beryllium greater than the ionisation energy of boron.

A

There is a slight decrease in IE1 between beryllium and boron as the fifth electron in boron is in the 2p subshell, which is further away from the nucleus than the 2s subshell of beryllium
Beryllium has a first ionisation energy of 900 kJ mol-1 as its electron configuration is 1s2 2s2
Boron has a first ionisation energy of 800 kJ mol-1 as its electron configuration is 1s2 2s2 2px1

83
Q

Why is the ionisation energy of oxygen greater than the first ionisation energy of nitrogen?

A

There is a slight decrease in IE1 between nitrogen and oxygen due to spin-pair repulsion in the 2px orbital of oxygen
Nitrogen has a first ionisation energy of 1400 kJ mol-1 as its electron configuration is 1s2 2s2 2px1 2py1 2pz1
Oxygen has a first ionisation energy of 1310 kJ mol-1 as its electron configuration is 1s2 2s2 2px2 2py1 2pz1

84
Q

Whys is there a great difference in successive ionisation energies?

A

A Level Chemistry-CIE (13 decks)

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