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AS Chemistry > Atomic structure > Flashcards

Atomic structure Flashcards

1
Q

What is the structure of an atom

A
  • Atoms are mostly made up of empty space around a very small, dense nucleus that contains protons and neutrons
  • Negatively charged electrons are found in orbitals in the empty space around the nucleus
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2
Q

What is the charge of the nucleus

A

The nucleus has an overall positive charge

The protons have a positive charge and the neutrons have a neutral charge

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3
Q

What are subatomic particles

A

The protons, neutrons and electrons that an atom is made up of are called subatomic particles

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4
Q

what is the relative charge of a proton

A

+1

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5
Q

what is the relative mass of a proton

A

1

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6
Q

what is the relative charge of a neutron

A

0

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7
Q

what is the relative mass of a neutron

A

1

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8
Q

what is the relative charge of an electron

A

-1

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9
Q

what is the relative mass of an electron

A

1/1836

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10
Q

what is the atomic/proton number

A
  • the number of protons in the nucleus of an atom and has the symbol Z
  • The atomic number is also equal to the number of electrons present in a neutral atom of an element
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11
Q

what is the mass/nucleon number

A

the total number of protons + neutrons in the nucleus of an atom, and has the symbol A

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12
Q

how to calculate the number of neutrons

A

Number of neutrons = mass number – atomic number

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13
Q

what are nucleons

A

Protons and neutrons are also called nucleons, because they are found in the nucleus

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14
Q

what is the mass distribution of an atom

A
  • The mass of an atom is concentrated in the nucleus, because the nucleus contains the heaviest subatomic particles (the neutrons and protons)
  • The mass of the electron is negligible
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15
Q

what is the charge distribution of an atom

A
  • The nucleus is also positively charged due to the protons
  • Electrons orbit the nucleus of the atom, contributing very little to its overall mass, but creating a ‘cloud’ of negative charge
  • The electrostatic attraction between the positive nucleus and negatively charged electrons orbiting around it is what holds an atom together
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16
Q

How do electrons behave in an electric field (in-between electrically charged plates)

A
  • When a beam of electrons is fired past the electrically charged plates, the electrons are deflected very easily away from the negative plate towards the positive plate
  • This proves that the electrons are negatively charged; like charges repel each other
  • It also shows that electrons have a very small mass, as they are easily deflected
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17
Q

How do protons behave in an electric field (in-between electrically charged plates)

A
  • A beam of protons is deflected away from the positive plate and towards the negative plate
  • This proves that the proton is positively charged
  • As protons are deflected less than electrons, this also shows that protons are heavier than electrons
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18
Q

How do neutrons behave in an electric field (in-between electrically charged plates)

A
  • A beam of neutrons is not deflected at all
  • Which proves that the particle is neutral in character; it is not attracted to, or repelled by, the negative or positive plate
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19
Q

How are ions formed

A

Ions on the other hand are formed when atoms either gain or lose electrons, causing them to become charged

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20
Q

What is the atomic radius

A

it is half the distance between two nuclei of two covalently bonded atoms of the same type

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21
Q

What trends do atomic radii show across the periodic table

A
  • They generally decrease across each Period

- They generally increase down each Group

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22
Q

Why do atomic radii generally decrease across each period

A
  • Atomic radii decrease as you move across a Period as the atomic number increases (increased positive nuclear charge) but at the same time extra electrons are added to the same principal quantum shell
  • The larger the nuclear charge, the greater the pull of the nuclei on the electrons which results in smaller atoms
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23
Q

Why do atomic radii generally increase moving down a group

A
  • Atomic radii increase moving down a Group as there is an increased number of shells going down the Group
  • The electrons in the inner shells repel the electrons in the outermost shells, shielding them from the positive nuclear charge
  • This weakens the pull of the nuclei on the electrons resulting in larger atoms
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24
Q

what is ionic radius

A

The ionic radius of an element is a measure of the size of an ion

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25
Q

what trends do ionic radii show across the periodic table

A
  • Ionic radii increase with increasing negative charge

- Ionic radii decrease with increasing positive charge

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26
Q

Why do ionic radii increase with increasing negative charge

A

Ions with negative charges are formed by atoms accepting extra electrons while the nuclear charge remains the same
The outermost electrons are further away from the positively charged nucleus and are therefore held only weakly to the nucleus which increases the ionic radius

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27
Q

Why do ionic radii decrease with increasing positive charge

A

Positively charged ions are formed by atoms losing electrons
The nuclear charge remains the same but there are now fewer electrons which undergo a greater electrostatic force of attraction to the nucleus which decreases the ionic radius

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28
Q

What are isotopes

A

Isotopes are atoms of the same element that contain the same number of protons and electrons but a different number of neutrons

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29
Q

How do isotopes differ from each other in terms of properties

A

Isotopes have similar chemical properties but different physical properties

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30
Q

why do isotopes have similar chemical properties

A
  • they have the same number of electrons in their outer shells
  • Electrons take part in chemical reactions and therefore determine the chemistry of an atom
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31
Q

why do isotopes have different physical properties

A
  • The only difference between isotopes is the number of neutrons
  • Since these are neutral subatomic particles, they only add mass to the atom
  • As a result of this, isotopes have different physical properties such as small differences in their mass and density
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32
Q

what is electronic configuration

A
  • the arrangement of electrons in an atom is called the electronic configuration
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33
Q

what are principal energy levels/principal quantum shells

A

what electrons are arranged around the nucleus in

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34
Q

what are principal quantum numbers (n)

A

they are used to number the energy levels or quantum shells as well as the energy of the electrons in that shell

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35
Q

how many electrons can n = 1 (principle quantum shell/energy level 1) hold

A

up to 2 electrons

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36
Q

how many electrons can n = 2 (principle quantum shell/energy level 2) hold

A

up to 8 electrons

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37
Q

how many electrons can n = 3(principle quantum shell/energy level 3) hold

A

up to 18 electrons

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38
Q

how many electrons can n = 4 (principle quantum shell/energy level 4) hold

A

up to 32 electrons

39
Q

what are subshells

A
  • the principal quantum shells are split into subshells

- there are a maximum of 4 subshells (s, p, d and f)

40
Q

how many electrons can subshell s hold

A

2 electrons max

41
Q

how many electrons can subshell p hold

A

6 electrons max

42
Q

how many subshells can subshell d hold

A

10 electrons max

43
Q

how many subshells can subshell f hold

A

14 electrons max

44
Q

what subshells do n = 1 have?

A

subshell s

45
Q

what subshells do n = 2 have

A

subshell s and p

46
Q

what subshells do n = 3 have

A

subshells s, p and d

47
Q

what subshells do n > 3 have?

A

subshells s, p, d and f

48
Q

what happens to the order of subshells for the higher principle quantum shells

A

they overlap

49
Q

what are orbitals

A
  • orbitals are the movement pattern of electrons
  • subshells contain one or more atomic orbitals
  • orbitals exist at specific energy levels and electrons can only be found at these specific levels, not inbetween them
  • each atomic orbital can be occupied by a maximum of 2 electrons
50
Q

how many orbitals does subshell S have

A

1 orbital

51
Q

how many orbitals does subshell P have

A

3 orbitals

52
Q

how many orbitals does subshell D have

A

5 orbitals

53
Q

how many orbitals does subshell F have

A

7 orbitals

54
Q

what are the shapes of the s orbitals

A

https://cdn.savemyexams.co.uk/wp-content/uploads/2020/11/1.1-Atomic-Structure-Orbitals.png

  • The s orbitals are spherical in shape
  • The size of the s orbitals increases with increasing shell number
55
Q

what are the shapes of the p orbitals

A

https://cdn.savemyexams.co.uk/wp-content/uploads/2020/11/1.1-Atomic-Structure-Orbitals.png

  • The p orbitals are dumbbell-shaped
  • Every shell has three p orbitals except for the first one (n = 1)
  • The p orbitals occupy the x, y and z-axis and point at right angles to each other so are oriented perpendicular to one another
  • The lobes of the p orbitals become larger and longer with increasing shell number
56
Q

What is the ground state

A

the ground state is the most stable electronic configuration of an atom which has the lowest amount of energy

57
Q

how is the ground state achieved

A

by filling the subshells of energy with the lowest energy first (1s)

58
Q

what is the pattern of the order of subshells in terms of increasing energy

A

The order of the subshells in terms of increasing energy does not follow a regular pattern at n= 3 and higher

59
Q

what are the differences in energy levels between the orbitals in a subshell

A

In the ground state, orbitals in the same subshell have the same energy and are said to be degenerate, so the energy of a Px orbital is the same as a Py orbital

60
Q

What is electronic configuration?

A

The electronic configuration gives information about the number of electrons in each shell, sub shell and orbital of an atom

61
Q

How are the subshells filled?

A

The subshells are filled in order of increasing energy

62
Q

What are electron spins?

A

Electrons can be imagined as small spinning charges which rotate around their own axis in either a clockwise or anti clockwise direction
- the spin of the electron is represented by its direction

63
Q

What is spin-pair repulsion

A

Electrons with similar spin repel each other which is also called spin-pair repulsion

64
Q

How do electrons occupy orbital within subshells based on spin charge ?

A
  • Electrons will occupy separate orbitals in the same subshell to minimizethis repulsion and have their spin in the same direction
  • electrons are only paired when there are more empty orbitals available within a subshell in which case the spins are the opposite spins to minimize repulsion
65
Q

Why do negatively charged electrons occupy the same region of space in orbitals?

A

Because the energy required to jump successive empty orbital is greater than the inter-electron repulsion
- for this reason they pair up and occupy the lower energy levels first

66
Q

What is electron box notation?

A
  • electronic configuration can be represented using the electrons in boxes notation
  • each box represents an atomic orbital
  • the boxes are arranged in order of increasing energy from lowest to highest
  • the electrons are represented by opposite arrows to show the spin of the electrons
67
Q

What is a free radical?

A
  • a free radical is a species with one or more unpaired electron
  • the unpaired electron in the free radical is shown as a dot
68
Q

What are the electronic configuration exceptions?

A
  • Cr is [Ar] 3d5 4s1 not [Ar] 3d4 4s2
  • Cu is [Ar] 3d10 4s1 not [Ar] 3d9 4s2
  • this is because the configurations are energetically stable
69
Q

What is the full order of filling of the orbitals?

A

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

70
Q

how is the periodc table organized based on orbital electronic configurations

A
  • s block element (group 1-2) have their valence electrons in an s orbital
  • p block elements (group 13-18) have their valence electrons in a p orbital
  • d block elements (group 3-12) have their valence electrons in a d orbital
  • f block elements (rare+radioactive) have their valence electrons in an f orbital
71
Q

what are ionisation energies

A

the ionisation energy (IE) of an element is the amount of energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous ions

72
Q

how are ionisation energies measured

A

ionisation energies are measured under standard conditions which re 298K and 1 atm

73
Q

what are the units of ionisation energies

A

kilojoules per mole (kJ mol^-1)

74
Q

what is meant by successive ionisation energies

A
  • the electrons from an atom can be continued to be removed until only the nucleus is left
  • this sequence of ioniation energies is called successive ionisation energies
75
Q

what is the trend of ionisation energies across the periodic table

A

the group 1 metals have a relatively low ionisation energy where as the noble gases have very high ionisation energies

76
Q

what are the four factors affecting ionization energy

A
  • size of nuclear charge
  • distance of outer electrons from the nucleus
  • shielding effect of inner electrons
  • spin-pair repulsion
77
Q

how does size of nuclear charge affect ionization energy

A

the nuclear charge increases with increasing atomic number , which means that there are greater attractive forces between the nucleus and electrons, so more energy is required to overcome these attractive forces when removing an electron

78
Q

how does distance of outer electrons from the nucleus affect ionization energy

A

electrons in shells that are further away from the nucleus are less attracted to the nucleus (the nuclear attraction is weaker) so the further the outer electron shell is from the nucleus, the lower the ionisation energy

79
Q

how does shielding effect of inner electrons affect ionization energy

A

the shielding effect is when the electrons in full inner shells repel electrons in outer shells, preventing them from feeling the full nuclear charge, so the more shells an atom has, the greater the shielding effect, and the lower the ionisation energy

80
Q

how does spin-pair repulsion affect ionization energy

A

-electrons in the same atomic orbital in a subshell repel each other more than electrons in different atomic orbitals which make it easier to remove and electron (which is why the first ionization energy is always the easiest

81
Q

what is the trend of ionization energies across a period

A

the ionization energy over a period increases

82
Q

why does the ionisation energy over a period increase?

A
  • Across a period the nuclear charge increases so the atomic radius of the atoms to decrease, as the outer shell is pulled closer to the nucleus, so the distance between the nucleus and the outer electrons decreases
  • The shielding by inner shell electrons remain reasonably constant as electrons are being added to the same shell
83
Q

what is the trend in ionisation energies between periods

A

There is a rapid decrease in ionisation energy between the last element in one period, and the first element in the next period

84
Q

why is there there is a rapid decrease in ionisation energy between the last element in one period, and the first element in the next period

A
  • There is increased distance between the nucleus and the outer electrons as you have added a new shell
  • There is increased shielding by inner electrons because of the added shell
  • These two factors outweigh the increased nuclear charge
85
Q

what are the exceptions in the trend in ionisation energies across a period

A
  • There is a slight decrease in IE1 between beryllium and boron
  • There is a slight decrease in IE1 between nitrogen and oxygen and phosphorus
86
Q

why is there a slight decrease in first ionization energy between beryllium and boron

A
  • 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
87
Q

why is there a slight decrease in first ionization energy between nitrogen and oxygen and phosphorus

A
  • 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
88
Q

what is the trend of ionization energies down a group

A

the ionisation energy down a group decreases

89
Q

why does the ionization energy decrease down a group

A
  • the distance between the nucleus and outer electron increases as you descend the group as more shells are added
  • The shielding by inner shell electrons increases as there are more shells of electrons
  • These factors outweigh the increased nuclear charge, meaning it becomes easier to remove the outer electron as you descend a group
90
Q

what are the ionization energy trends across a period

A
  • increase in nuclear charge
  • shell number is the same and distance of outer electron to nucleus is the same
  • shielding remains reasonably constant
  • decreased atomic/ionic radius
  • the outer electron is held more tightly to the nucleus so it gets harder to remove it
91
Q

what are the ionization energy trends down a group

A
  • increase in nuclear charge
  • increase in shells
  • distance of outer electron to nucleus increases
  • shielding effect increases, therefore, the attraction of valence electrons to the nucleus decreases
  • increased atomic/ionic radius
  • the outer electron is held more loosely to the nucleus so it gets easier to remove
92
Q

why do the sucessive ionisation energies of an element increase?

A
  • because once you hve removed the outer electron from an atom, you have formed a positive ion
  • removing an electron from a positive ion is more difficult than from a neutral atom
  • as more electrons are removed, the attractive forces increase due to decreasing shieldingand an increase in the proton:electron ratio
  • the increase in ionization energy is not constant and depends on the atoms electronic configuration
93
Q

how can the electronic configuration of an atom be determined by their ionization energies

A

By analyzing where the large jumps appear and the number of electrons removed when these large jumps occur, the electron configuration of an atom can be determined

AS Chemistry (22 decks)

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