Atomic Structure Flashcards

1
Q

Define nucleons

A

The nucleus is made up of particles called nucleons. Protons and neutrons are also called nucleons.

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2
Q

What is the mass, charge and location of the particle electrons.

A

Mass: 1/1836
Charge: -1
Location: Orbitals

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3
Q

What is the mass, charge and location of the particle protons.

A

Mass=1
Charge=+1
Location= nucleus

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4
Q

What is the mass, charge and location of the particle neutrons

A

Mass=1
Charge=0
Location=Nucleus

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5
Q

Describe the behaviour of a beam of electrons when fired past 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.

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6
Q

Explain the behaviour of a beam of electrons when fired past electrically charged plates.

A

Their behaviour proves that the electrons are negatively charged as like charges repel each other.
This also proves that electrons have a very samll mass, as they are easily deflected.

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7
Q

Describe the behaviour of a beam of protons when fired past the electrically charged plates.

A

A beam of protons when fired past electrically charged plates are deflected away from the positive plate towards the positive plate.

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8
Q

Explain the behaviour of a beam of protons when fired past electrically charged plates.

A

The behaviour of a beam of protons when fired past electrically charged plates proves that the proton is positively charged.
As proton are deflected less than electrons, this shows that protons are heavier than electrons,

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9
Q

Describe the behaviour of a beam of neutrons when fired past electrically charged plates.

A

A beam of neutrons is not deflected at all.

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10
Q

Explain the behaviour of a beam of neutrons when fired past electrically charged plates.

A

The behaviour of a beam of neutrons proves that the particle is neutral in character; it is not attracted or repelled by the negative or positive plate.

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11
Q

Define atomic number.

A

The atomic number (or proton number) is the number of protons in the nucleus of an atom and has the symbol z.

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12
Q

Define mass number.

A

The mass number (or nucleon number) is the total number of protons+neutrons in the nucleus of an atom, and has the symbol A.

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13
Q

How can the number of neutrons be calculated?

A

The number of neutrons can be calculated:

Number of neutrons=mass number-atomic number.

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14
Q

How can you read nuclide notation?

A

CHECK NOTES

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15
Q

Define an ion and how it is formed.

A

An ion is a charged atom. Ions are formed when atoms either gain or lose electrons, causing them to be charged.

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16
Q

What are the two types of ions?

A

There are two types of ions:
Cations
Anions

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17
Q

Define a cation.

A

A cation is an ionic species with a positive charge.

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18
Q

Define an anion.

A

An anion is an ionic species with a negative charge.

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19
Q

Define a free radical.

A

Free radicals are highly reactive atoms that have one unpaired electron.

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20
Q

Are free radicals stable or unstable. Explain why.

A

Free radicals are unstable as there is tendency for unpaired electrons to pair up and so the free radicals react very quickly.

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21
Q

How are free radicals formed?

A

Free radicals can be formed when a covalent bond is broken by energy supplied by UV light.

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22
Q

Define 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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23
Q

What is the symbol for an isotope?

A

The symbol for an isotope is the chemical symbol (or word) followed by a dash and then the mass number.
For e.g. carbon-12 and carbon-14

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24
Q

Do isotopes have similar or different chemical properties and why?

A

Isotopes have similar chemical properties. This is because they have the same number of electrons in their outer shell. Electrons take part in chemical reaction therefore determines the chemistry of an atom.

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25
Q

Do isotopes have similar or different physical properties and why?

A

isotopes of the same element display different physical properties such as small differences in their mass and density as they have different number of neutrons.

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26
Q

What are the three isotopes of hydrogen naturally present?

A

There are three isotopes of hydrogen naturally present.

  1. Protium (Proton=1 Neutron=0)
  2. Deuterium ( Proton=1 Neutron=1)
    3.Tritium( Proton=1 neutron=2)

Protium or common hydrogen is the only atom without a neutron in its nucleus.

27
Q

Define Principle Quantum Number.

A

Electrons are arranged around the nucleus in principle quantum shells. Principle Quantum numbers (n) are used to number quantum shells and give information about the energy of the shell. The higher the principle quantum number, the greater the energy of the shell. The lower the principle quantum number, the closer the shell is to the nucleus. Each principle quantum number has a fixed number of electrons it can hold.

28
Q

What formula is used to determine how many electrons are in each quantum shell?

A

2n^” where (n) represents the shell number.
1=2
2=8
3=18
4=32

29
Q

Define subshells:

A

The principle quantum shells are split into subshells which are given the letters s, p ,d and f. The energy of the electrons in the subshells increases in the order

s<p<d<f

30
Q

Define orbitals:

A

The region of space around the nucleus where the probability of finding an electron is maximum is called the orbital.

31
Q

What is the fixed number of orbitals in and electrons in each subshell?

A

Each subshell contains a fixed number of orbitals that contains electrons.

An atomic orbital can only be occupied by one or two electrons only.

S subshell=max of 2 electrons=1 orbital
P subshell=max of 6 electrons=3 orbitals
d subshell=max of 10 electrons=5 orbitals
f=subshell=max of 14 electrons= 7 orbitals

32
Q

Describe the shape of the s orbital.

A

The s orbital is spherically shaped.(check diagram)

33
Q

Describe the shape of the p orbital.

A

The p orbital is dumbbell shaped. (check diagram)

34
Q

What happens to the three dimensional shape of an orbital as the principle quantum number increases.

A

The size of the orbital increases with increasing principle quantum number as the energy increases.

35
Q

Why are the Px, Py and Pz orbitals called degenerate orbitals?

A

Px, Py, Pz orbitals are called degenerate orbitals because they have the same energy.

36
Q

Define the azumuthal quantum number.

A

The azumthal quantum number (l) gives information about the energy of the subshell.

L=0→s
L=1→p
L=2→d
L=3→f

The higher the value of (l) the greater the energy of the subshell.

37
Q

Define electron configuration.

A

The electron configuration tells us how the electrons in an atom or ion are arranged in their shells.

38
Q

What are the principles upon which the filling of electrons is based upon?

A
  1. Aufbau Principle
  2. Pauli’s exclusion principle
  3. Hund’s rule
  4. Stability of atomic orbitals
39
Q

Explain the Aufbau principle.

A

Electrons always fell the lowest energy subshell first.

If two different orbitals have identical energies, the electrons fill first in that whose value of (n) is less.

40
Q

Explain Pauli’s exclusion principle.

A

No orbital can accommodate more than two electrons. If there are two electrons in an orbital, they must have opposite spins.

Electrons don’t experience repulsion because of their spin movement which is in opposite direction.

41
Q

Explain Hund’s Rule.

A

When their are a number of orbitals of equal energy, electrons first fill them up individually and then get paired. By filling up individually, mutual repulsion between electrons is avoided and thereby maximum stability is achieved.

42
Q

Explain Stability of atomic orbitals.

A

(this rule is only applicable to chromium and copper)

Completely filled atomic orbital is more stable than half filled orbital which in turn is more stable than incomplete or partially filled orbital.

43
Q

What is the formula for electronic configuration.

A

1s^2

1=principle quantum number
s=orbital
^2=electron

44
Q

Which elements are called s, p, d and f block elements?

A

The s block elements have their valence electron (s) in the s orbital.
The p block element have their valence electron (s) in the p orbital.
The d block elements have their valence electron (s) in the d orbital.
The F block element have their valence electron (s) in the f orbital.

45
Q

Define ionisation energy.

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.

46
Q

What are the standard conditions in which ionisation energy is measured in?

A

Ionisation energies are measured under standard conditions which are 298K and 101Kpa.

47
Q

What is the unit in which ionisation energy is measured in?

A

The units of IE are kilojoule’s per mole (KJ mol-1)

48
Q

Define first ionisation energy.

A

The fist ionisation energy is the energy required to remove one mole of electrons from one mole of atoms of an element to form one mole of gaseous ions.

For e.g.

Ca(g)→ca+(g) +e-

49
Q

Define the second ionisation energy.

A

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

50
Q

Is Ionisation energy endothermic or exothermic?

A

Ionisation energy is an endothermic process, because energy is required to break the force of attraction between the electron and the central positive nucleus.

51
Q

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

A

The second ionisation energy is greater than the first ionisation energy as there is a stronger attraction between more proton and less electrons in the 1+ ion.

52
Q

What are the factors which impact ionisation energy?

A

The size of the first ionisation energy is affected by four factors:
* Size of the nuclear charge
* Distance of outer electrons from the nucleus
* Shielding effect of inner electrons
* Spin-Pair repulsion

53
Q

What is the trend in ionisation energy across a period?

A

First Ionisation energy increases across a period.

54
Q

What is the trend in ionisation energy down a group?

A

First ionisation energy decreases down a group.

55
Q

Explain the trend in ionisation energy across a period?

A

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 electrons remains reasonably constant as the electrons are being added to the same shell. It becomes harder to remove an electron across a period so more ionisation energy is required.

56
Q

Explain the trend in ionisation energy from one period to the next.

A

The ionisation energy over a period increases due to the following factors: 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.

57
Q

What is the trend in ionisation energy from one period to the next?

A

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

58
Q

Explain the trend in ionisation energy down a group?

A

The nuclear charge increases but as the atomic radius of the atoms increases as you are adding more shells of electrons making shells bigger. The shielding by inner electrons also increases as you descend a group. These factors outweigh the increases nuclear charge, meaning it becomes easier to remove an electron as you descend a group.

59
Q

what is meant by successive ionisation energies?

A

The removal of electrons from an atom one by one successively is called successive ionisation energies.

60
Q

What is the trend in successive ionisation energies?

A

The successive ionisation energies of an element increase, because once you have removed an electron from an atom you have formed a positive ion. This makes removing an electron more difficult as more electrons are removed the attractive forces increase due to decreased shielding and an increase in the proton to electron ratio.

61
Q

Define ground state.

A

The ground state is the most stable electronic configuration of an atom which has the lowest amount of energy. This is acheived by filling the lowest energy orbital first.

62
Q

Define excited state:

A

When the electrons absorb energy and jump to outer orbits this state is called excited state.

63
Q

Define isoelectronic.

A

Isoelectronic means shaving the same number of electrons or the same electronic structure.