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AS Chemistry 9701 > Atomic Structure > Flashcards

Atomic Structure Flashcards

1
Q

What is the structure of an Atom

A

Atoms are mostly empty space surrounding a very small, dense nucleus that contains
protons and neutrons; electrons are found in shells in the empty space around the nucleus

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2
Q

What is the overall charge of the nucleus

A

+

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3
Q

Where are negatively charged electrons found

A

Orbitals in the empty space surrounding the nucleus

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4
Q

What are protons, electrons and neutrons

A

Subatomic Particles

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5
Q

How are subatomic particles’ charges and masses calculated

A

They are compared to each other so are
1. Relative Atomic Mass
2. Relative Atomic Charge

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6
Q

What are relative charges of subatomic particles

A

Proton = +1
Neutron = 0
Electron = -1

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7
Q

What are relative masses of subatomic particles

A

Proton = 1
Neutron = 1
Electron = 1/1836

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8
Q

What is the charge of a proton and electron

A

Proton = +1.602 x 10^-19
Electron = +1.602 x 10^-19

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9
Q

What is the mass/nucleon number

A

Number of protons and neutrons in the nucleus and the mass of the atom

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10
Q

What is the proton/atomic number

A

Number of protons or electrons in the atom

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11
Q

What is the distribution of mass in an atom

A

All the mass is concentrated in the nucleus

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12
Q

What is the distribution of charge in an atom

A

The nucleus is positively charged due to protons

Electrons orbit the nucleus and create a ‘cloud’ of negative charge

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13
Q

What holds an atom together

A

The electrostatic attraction between positive nucleus and negative charge of orbiting electrons

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14
Q

How do subatomic particles behave at the same velocity in a electric field

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

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

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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15
Q

What is an atom and ion

A

An atom is neutral and has no overall charge

An ion is formed when atoms gain or lose electrons

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16
Q

What is the atomic radius

A

The measure of the size of an atom.
It is half of the distance between nuclei of 2 covalently bonded atoms of the same type.

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17
Q

What are the variations of atomic radii

A

Decrease down each period
Increase down each group

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18
Q

Why do atomic radii decrease down a period

A

Electron Shell Theory. 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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19
Q

Why do atomic radii increase 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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20
Q

What is the ionic radius

A

Is the measure of the size of an ion - distance from the nucleus of the ion to the valence shell

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21
Q

What are the variations of ionic radii

A

Increase with increasing negative charge
Decrease with increasing positive charge

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22
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

The greater the negative charge, the larger the ionic radius

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23
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

The greater the positive charger, the smaller the ionic radius

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24
Q

What is an isotope

A

Isotopes are atoms of the same element with same number of protons/proton number but different number of neutrons/nucleon number

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25
What is the notation of Isotopes
The name of an isotope is the chemical symbol (or word) followed by a dash and then the mass number Eg. carbon-12 and carbon-14 are isotopes of carbon containing 6 and 8 neutrons respectively Or a large letter with 2 small letters to the left of it. Top is mass number and bottom is proton number
26
Why do isotopes of the same element have the same chemical properties
Isotopes of the same element display the same chemical characteristics This is because 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
27
Why do isotopes of the same element have different physical properties
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
28
What is the arrangement of electrons called
The Electronic configuration
29
What are electrons arranged in
Principal energy levels of Principal quantum shells
30
What are principal quantum numbers (n)
Higher numbers, higher the energy of the shell and further it is from nucleus Lower numbers, closer the shell is to the nucleus and less energy it holds They number the principal quantum shells They describe the energy of an electron in a orbital.
31
How many electrons can each (n) hold
n = 1 holds up to 2 n = 2 holds up to 8 n = 3 holds up to 18 n = 4 holds up to 32
32
What are sub-shells
They are subdivisions of a shell or principal quantum shell and there are 4 s
33
How does an atom get an F subshell
Having more than 57 electrons
34
List the subshells in increasing energy
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 4f
35
What are orbitals
1 or more orbitals exist within a subshell. Orbitals exist at specific energy levels and electrons can only be found at these specific levels, not in between them.
36
How many max electrons in one orbital
2
37
How many orbitals in a subshell
s = 1 = 2 electrons p = 3 d = 5 f = 7 = 14 electrons
38
What shape do orbitals have
Specific 3d shapes Savemyexams image
39
What shape do S and P orbitals have
S orbitals are spherical in shape The size of the S orbital increase with increasing shell number. Ex. n=3 (s) is bigger than n=1 (s) P orbitals are dumbbell shaped. 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
What does a complete shell look like
Savemyexams
41
What is the ground state and how is it achieved including anomalies in subshell energies
The ground state is the most stable electronic configuration of an atom which has the lowest amount of energy This is achieved by filling the sub-shells with the lowest energy first (1s) The order of the sub-shells in terms of increasing energy does not follow a regular pattern at n = 3 and higher Savemyexams Image
42
How does energy of individual orbitals in the same subshell in ground state compare
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
43
What is degenerate
When all the orbitals in the same subshell have the same energy they are 'Degenerate'
44
What is the electronic configuration
The electron configuration gives information about the number of electrons in each shell, sub-shell and orbital of an atom The sub-shells are filled in order of increasing energy to maintain balance of inter-electron repulsion and attraction of electrons to nucleus when energy of electrons increase.
45
Representation of electronic configurations
Savemyexams image
46
How do electrons move
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. When North is up = clockwise. Vice versa. Savemyexams image
47
What is spin-pair repulsion
Electrons with similar spin repel each other which is also called spin-pair repulsion
48
How do atoms minimize the spin pair repulsion
Electrons will therefore occupy separate orbitals in the same sub-shell to minimise this repulsion and have their spin in the same direction
49
When are electrons paired
Electrons are only paired when there are no more empty orbitals available within a sub-shell in which case the spins are the opposite spins to minimise repulsion
50
What is inter-electron repulsion
Repulsion between negatively charged electrons
51
Why don't electrons fill up other orbitals on other subshells than pair up when there's inter-electron repulsion
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
52
What does electron box configuration show and do
The electron configuration can also 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
53
What is the rule for 3D
Always fill up 4S first 3D cannot be full or half full
54
What is a free radical and how is it formed
A free radical is a species with one or more unpaired electron The unpaired electron in the free radical is shown as a dot. If Cl radical = Cl ' where dot is in middle height of letter 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.
55
What are the full and shorthand electronic configuration
The full electron configuration describes the arrangement of all electrons from the 1s sub-shell up The shorthand electron configuration includes using the symbol of the nearest preceding noble gas to account for however many electrons are in that noble gas
56
How do ions gain or lose electrons in subshells including 3d and 4s
Negative ions are formed by adding electrons to the outer sub-shell Positive ions are formed by removing electrons from the outer sub-shell The transition metals fill the 4s sub-shell before the 3d sub-shell but lose electrons from the 4s first and not from the 3d sub-shell
57
Full electronic configuration representation of K
Potassium - 19 electrons 1s^2 2s^2 2p^6 ... 1s, 1 represents energy number s^2, 2 represents number of electrons in subshell
58
Shorthand electronic configuration representation of K
Potassium - 19 electrons [Ar] 4s^1 [Ar] is Argon - the nearest noble gas to potassium with 18 electrons so we can ignore the 18 electrons. 19-18 = 1 So last electron is in 4s^1
59
What is ionisation
Ionisation is the process by which an electron is removed from an atom or a molecule
60
What is ionisation energy
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 Ionisation energies are measured under standard conditions which are 298 K and 101 kPa The units of IE are kilojoules per mole (kJ mol-1) The values for ionisation energies are always positive as this is an endothermic process This is because energy is required to break the force of attraction between the electron and the central positive nucleus
61
What is first ionisation energy
The first ionisation energy (IE1) is the energy required to remove one mole of electrons from one mole of atoms of an element to form one mole of 1+ ions Ex. Ca(g) ---> Ca^+ (g) + e^-1 , IE1 = 590
62
What is second ionisation energy
The second ionisation energy (IE2) of an element is the amount of energy required to remove one mole of electrons from one mole of gaseous ions of an element to form one mole of gaseous 2+ ions E.g. the second ionisation energy of gaseous calcium: Ca+ (g) → Ca2+ (g) + e- IE2 = +1145.4 kJ mol-1 IT STARTS FROM +1 NOT NEUTRAL
63
How is the first ionisation energy affected by size of nuclear charge
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
64
How is the first ionisation energy affected by distance of outer electrons to nucleus
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
65
How is the first ionisation energy affected by shielding effect of inner electrons
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
66
How is the first ionisation energy affected by spin-pair repulsion
Electrons in the same atomic orbital in a subshell repel each other more than electrons in different atomic orbitals which makes it easier to remove an electron (which is why the first ionisation energy is always the lowest)
67
How does FIRST ionisation energy change in period table
Increases across period Decreases down group
68
Why does first ionisation energy increase over a period
Across a period the nuclear charge increases This causes 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 It becomes harder to remove an electron as you move across a period; more energy is needed So, the ionisation energy increases
69
Why is there a rapid decrease in IE1 between last element of 1st period and first element of 2nd period
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
70
Why does ionisation energy decrease down a group
The number of protons in the atom is increased, so the nuclear charge increases But, the atomic radius of the atoms increases as you add more shells of electrons, making the atoms bigger So, the distance between the nucleus and outer electron increases as you descend the group 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 So, the ionisation energy decreases
71
How do successive ionisation energies differ and why
Increase This is because once you have 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 shielding and an increase in the proton to electron ratio The increase in ionisation energy, however, is not constant and is dependent on the atom's electronic configuration
72
How does difficulty of removing electron change
It is easy to remove electrons from a full subshell as they undergo spin-pair repulsion. It gets more difficult to remove electrons from principal quantum shells that get closer to the nucleus as there is less shielding and an increase in attractive forces between the electrons and nuclear charge.
73
What is required to remove outer electron and why
Energy is required to remove an outer shell electron as this involves breaking the attractive forces between the electron and the positively charged nucleus
74
How does Nuclear charge affect magnitude of the ionisation energy
Positive nuclear charge increases with increasing number of protons The greater the positive charge, the greater the attractive forces between the outer electron(s) and the nucleus More energy is required to overcome these forces so ionisation energy increases with increasing nuclear charge
75
How does shielding affect magnitude of ionisation energy
Electrons repel each other and electrons occupying the inner shells repel electrons located in shells further outside the nucleus and prevent them from feeling the full effect of the nuclear charge The greater the shielding effect is, the weaker the attractive forces between the positive nucleus and the negatively charged electrons Less energy is required to overcome the weakened attractive forces so ionisation energy decreases with increasing shielding effects
76
How does atomic/ionic radii affect magnitude of ionisation energy
The larger the radius, the greater the distance between the nucleus and the outer shell electron(s) Increasing distance weakens the strength of the attractive forces Larger atoms/ions also result in greater shielding due to the presence of more inner electrons Less energy is required to remove the outer shell electron(s) so ionisation energy decreases with increasing atomic/ionic radius
77
How does spin-pair repulsion affect magnitude of ionisation energy
Spin pair repulsion occurs when the electron being removed is spin paired with another electron in the same orbital The proximity of the like charges of electrons in the orbital results in repulsion Less energy is required to remove one of the electrons so ionisation energy decreases when there is spin-pair repulsion
78
How does nuclear charge, shielding and atomic/ionic radii come together
Nuclear charge and shielding link to atomic radii
79
What is the electronic configuration of chromium
Cr is [Ar] 4s1 3d5 not [Ar] 4s2 3d4
80
What is the electronic configuration of Cu
Cu is [Ar] 4s1 3d10 not [Ar] 4s2 3d9
81
Why are the Cu and Cr configurations different
This is because the [Ar] 4s1 3d5 and [Ar] 4s1 3d10 configurations are energetically stable

AS Chemistry 9701 (1 decks)

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