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A Level Chemistry-CIE > Atoms, moles and stoichiometry > Flashcards

Atoms, moles and stoichiometry Flashcards

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1
Q

What is an anode?

A

The positive electrode (where oxidation occurs).

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2
Q

What is the relative mass of an electron?

A

1/1836

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3
Q

What is an atomic number?

A

The number of protons in the nucleus of an atom. Also called the proton number.

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4
Q

What is an isotope?

A

Atoms of the same element with different mass numbers.

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5
Q

What is the mass number?

A

The number of protons+neutrons in an atom. Also called the nucleon number.

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6
Q

Do isotopes of a particular element have similar chemical properties?

A

Isotopes of a particular element have the same chemical properties because they have the same number of electrons.

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7
Q

Do isotopes of a particular element have similar physical properties?

A

Isotopes of a particular element have slightly different physical properties, such as small differences in mass and density because they have a different number of neutrons.

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8
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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9
Q

Explain the behaviour of subatomic particles in an electrical behaviour field.

A

Beams of protons and electrons are deflected by electrical fields but neutrons are not.

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10
Q

What is electron configuration?

A

The arrangement of electrons in an atom is called its electronic structure or electronic configuration.

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11
Q

What are quantum shells?

A

Electrons are arranged outside the nucleus in energy levels or quantum shells. These principle energy levels or principle quantum shells (symbolled n) 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 and so on.

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12
Q

Which electrons have greater energy?

A

The electrons in quantum shells further away from the nucleus have more energy and are held less tightly to the nucleus.

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13
Q

How many maximum electrons can principle quantum shells hold?

A

Shell 1: 2 electrons
Shell 2: 8 electrons
Shell 3: 18 electrons
Shell 4: 32 electrons

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14
Q

What is the first ionization energy, IE1?

A

First Ionisation energy is the energy needed to remove one mole of electrons from 1 mole of atoms of an element in the gaseous state to form one mole of gaseous ions.

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15
Q

What is ionisations energies units?

A

Its units are KJ mol-1

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16
Q

What is successive ionisation energy?

A

We can continue to remove electrons from an atom until only the nucleus is left. We call this sequence of ionization energies successive ionisation energies.

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17
Q

What is second ionisation energy?

A

The second ionisation energy is the energy needed to remove 1 electron from each ion of an element in 1 mole of gaseous +1 ions to form 1 mole of gaseous ions with a +2 charge.

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18
Q

What is the trend in ionisation energy across a period and down a group?

A

The first ionisation energy increases across a period and decreases down a group.

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19
Q

Why does ionisation energy increase 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.

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20
Q

Why does ionisation energy decrease down a group?

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.

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21
Q

What is shielding?

A

Shielding is the ability of inner shell electrons to reduce the effect of the nuclear charge on outer shell electrons.

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22
Q

What are the four factors which influence ionisation energy?

A

Four factors that influence ionisation energy are:
The size of the nuclear charge
Distance of Outer electrons from the nucleus
Shielding effect of inner electrons
Spin Pair repulsion

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23
Q

What is the general relationship between ionization energy and nuclear charge?

A

In general, ionisation energy increases as the proton number increases.

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24
Q

What is the general relationship between the distance of the outer electron shell and ionisation energy?

A

In general, the further the outer electron is from the nucleus, the lower the ionisation energy.

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25
Q

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

A

In general ionisation energy is lower as the number of full outer electron shells between the outer electrons and the nucleus increases.

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26
Q

What is the general trend between spin pair repulsion and ionisation energy?

A

Increased repulsion makes it easier to remove an electron so ionization energy decreases.

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27
Q

What is spin pair repulsion and how does it affect ionisation energy?

A

Electrons in the same atomic orbital in a subshell repel each other more than electrons in different atomic orbitals. This increased repulsion makes it easier to remove an electron so first ionisation energy decreases.

28
Q

What is the size of the nuclear charge and how does this affect ionisation energy?

A

As the atomic number increases, the positive charge in the nucleus increases. The bigger the positive charge, the greater the attractive force between the nucleus and the electrons. So more energy is required to overcome these attractive forces if an electron is to be removed.

29
Q

What is the distance of outer electrons from the nucleus and its effect on ionisation energy?

A

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

30
Q

What is the shielding effect of inner electrons and how does this effect ionisation energy?

A

As all electrons are negatively charged they repel each other. Electrons in full inner shells repel electrons in outer shells. Full inner shells of electrons prevent the outer electrons from feeling the nuclear charge. This is called the shielding effect. The greater the shielding of outer electrons by the inner electron shells, the lower the attractive forces between the nucleus and the outer electrons.

31
Q

What are subshells (subsidiary quantum shells)?

A

Subshells (subsidiary quantum shells) are regions of the principle quantum shells where electrons exist in defined spaces associated with particular amounts of energy. These are named s, p ,d and f.

32
Q

What are principle quantum shells split into?

A

The principle quantum shells, apart from the first, are split into sub-shells (sub-levels). Each principle quantum shell contains a different number of sub-shells. The subshells are distinguished by the letters s, p, d and f.

33
Q

What are atomic orbitals?

A

Each subshell contains one or more atomic orbitals. Atomic orbitals are regions of space outside the nucleus that can be occupied by a maximum of two electrons. Orbitals are named s, p, d and f. They have different shapes.

34
Q

How many orbitals can each subshell hold?

A

s subshell : 1 orbital
p subshell : 3 orbitals labelled px, py and pz
d subshell : 5 orbitals
f subshell : 7 orbitals

35
Q

What is the maximum number of electrons that each subshell can hold?

A

Each orbital can hold a maximum number of 2 electrons so the maximum number of electrons in each subshell are as follows:
s : 1 x 2 = total of 2 electrons
p : 3 x 2 = total of 6 electrons
d : 5 x 2 = total of 10 electrons
f : 7 x 2 = total of 14 electrons

36
Q

What is the order of increasing subshell energy?

A

The subshells increase in energy as follows: s ＜ p ＜ d ＜ f
The only exception to these rules is the 3d orbital which has slightly higher energy than the 4s orbital
Because of this, the 4s orbital is filled before the 3d orbital.

37
Q

What is a degenerate orbital?

A

All the orbitals in the same subshell have the same energy and are said to be degenerate
E.g. px, py and pz are all equal in energy

38
Q

Describe the shape of the s orbital?

A

The s orbitals are spherical in shape
The size of the s orbitals increases with increasing shell number
E.g. the s orbital of the third quantum shell (n = 3) is bigger than the s orbital of the first quantum shell (n = 1).
The s orbitals become larger with increasing principal quantum number.

39
Q

Describe the shape of the P orbitals.

A

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

40
Q

How can you identify which electron configuration of an atom has the lowest amount of energy?

A

The most stable electronic configuration of an atom is the one that has the lowest amount of energy.

41
Q

Which two element do not follow the usual pattern for writing electronic configuration?

A

Copper and chromium.
Copper has electronic configuration [AR] 3d^5 4s^1. (instead of 3d^4 4s^2)
Chromium has electronic configuration [AR] 3d^10 4s^1. (instead of [Ar] 3d^9 4s^2).
This is because the electronic configuration is more stable.

42
Q

Which elements are called s-block?

A

Elements in groups 1 and 2 have outer electrons in an s-subshell. They are therefore called the s-block.

43
Q

Which elements are called the p Block?

A

Elements in groups 13 to 18 (apart from He) have outer electrons in a p-subshell. They are therefore together called the P block.

44
Q

Which elements are called the d block ?

A

Elements that add electrons to the d-subshell are called the d-block elements. Most of of these are transition metals.

45
Q

What is spin pair repulsion?

A

Electrons can be imagined as small spinning charges which rotate around their own axis in either a clockwise or anticlockwise direction
The spin of the electron is represented by its direction.
electrons with similar spin repel each other which is also called spin-pair repulsion.

46
Q

Why do electrons occupy separate orbitals?

A

Electrons will therefore occupy separate orbitals in the same subshell to minimize this repulsion and have their spin in the same direction
Eg. if there are three electrons in a p subshell, one electron will go into each px, py and pz orbital

47
Q

When are electrons in a orbital within a subshell paired up?

A

Electrons are only paired when there are no more empty orbitals available within a subshell in which case the spins are the opposite spins to minimize repulsion
Eg. if there are four electrons in a p subshell, one p orbital contains 2 electrons with opposite spin and two orbitals contain one electron only

48
Q

Why do electrons occupy the same orbital instead of jumping to successive empty orbitals?

A

Even though there is repulsion between negatively charged electrons (inter-electrons repulsion), they occupy the same region of space in orbitals
This is because the energy required to jump to successive empty orbital is greater than the inter-electron repulsion
For this reason, they pair up and occupy the lower energy levels first.

49
Q

What is a free radical?

A

Free radical is a species with one (or sometimes more than one) unpaired electron.

50
Q

What can a free radical be represented as?

A

The unpaired electron in the free radical is shown as a dot. For example CL.

51
Q

How are free radials formed?

A

Free radicals are formed when a molecule undergoes homolytic fission where the two electrons of a covalent bond are split evenly between the two atoms.

52
Q

What is meant by species in terms of chemistry?

A

The word species refers to different particles, such as atoms, ions, molecules, free radicles or electrons when we want to write in general terms or about more than one type of particle.

53
Q

Describe how transition metals lose electrons slightly differently than the rest of the periodic table?

A

The transition metals fill the 4s subshell before the 3d subshell but lose electrons from the 4s first and not from the 3d subshell (the 4s subshell is lower in energy)

54
Q

What is the definition of atomic radius?

A

The atomic radius of an element is a measure of the size of an atom. It is half the distance between the two nuclei of two covalently bonded atoms of the same type.

55
Q

What is the definition of ionic radius?

A

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

56
Q

What is the trend in atomic radius down each group?

A

The atomic radius increases down any group.

57
Q

What is the trend in atomic radius across a period?

A

The atomic radius decreases across any period.

58
Q

Explain the trend in atomic radius down each 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.

59
Q

Explain the trend in atomic radius across a 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.

60
Q

What is the trend in ionic radius with increasing negative charge?

A

Ionic radii increase with increasing negative charge.

61
Q

What is the trend in ionic radius with increasing positive charge?

A

Ionic radii decrease with increasing positive charge.

62
Q

Explain the trend ionic radius 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
The greater the negative charge, the larger the ionic radius.

63
Q

Explain the trend in ionic radius 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
The greater the positive charger, the smaller the ionic radius

64
Q

Explain the rapid decreases in ionisation energy between the last element in one period and the first element in one period and the first element in the next period.

A

The forces of attraction between the positive nucleus and the outer negative electron decreases because:
The distance between the nucleus and the outer electron increases
The shielding by inner shells increases
These factors outweigh the increased nuclear charge.

65
Q

Why is there a slight decrease in first ionisation energy between beryllium and boron?

A

There is less attraction between the fifth electron in boron and the nucleus because:
The distance between the nucleus and the outer electron increases slightly.
The shielding by inner shells increases slightly
These two factors outweigh the increased nuclear charge.

A Level Chemistry-CIE (14 decks)

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