Atomic Structure

Question 1

Define nucleons

Answer 1

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

Question 2

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

Answer 2

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

Question 3

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

Answer 3

Mass=1
Charge=+1
Location= nucleus

Question 4

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

Answer 4

Mass=1
Charge=0
Location=Nucleus

Question 5

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

Answer 5

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.

Question 6

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

Answer 6

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.

Question 7

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

Answer 7

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

Question 8

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

Answer 8

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,

Question 9

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

Answer 9

A beam of neutrons is not deflected at all.

Question 10

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

Answer 10

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.

Question 11

Define atomic number.

Answer 11

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

Question 12

Define mass number.

Answer 12

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

Question 13

How can the number of neutrons be calculated?

Answer 13

The number of neutrons can be calculated:

Number of neutrons=mass number-atomic number.

Question 14

How can you read nuclide notation?

Answer 14

CHECK NOTES

Question 15

Define an ion and how it is formed.

Answer 15

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

Question 16

What are the two types of ions?

Answer 16

There are two types of ions:
Cations
Anions

Question 17

Define a cation.

Answer 17

A cation is an ionic species with a positive charge.

Question 18

Define an anion.

Answer 18

An anion is an ionic species with a negative charge.

Question 19

Define a free radical.

Answer 19

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

Question 20

Are free radicals stable or unstable. Explain why.

Answer 20

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

Question 21

How are free radicals formed?

Answer 21

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

Question 22

Define Isotopes.

Answer 22

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

Question 23

What is the symbol for an isotope?

Answer 23

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

Question 24

Do isotopes have similar or different chemical properties and why?

Answer 24

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.

Question 25

Do isotopes have similar or different physical properties and why?

Answer 25

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.

Question 26

What are the three isotopes of hydrogen naturally present?

Answer 26

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.

Question 27

Define Principle Quantum Number.

Answer 27

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.

Question 28

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

Answer 28

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

Question 29

Define subshells:

Answer 29

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

Question 30

Define orbitals:

Answer 30

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

Question 31

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

Answer 31

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

Question 32

Describe the shape of the s orbital.

Answer 32

The s orbital is spherically shaped.(check diagram)

Question 33

Describe the shape of the p orbital.

Answer 33

The p orbital is dumbbell shaped. (check diagram)

Question 34

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

Answer 34

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

Question 35

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

Answer 35

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

Question 36

Define the azumuthal quantum number.

Answer 36

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.

Question 37

Define electron configuration.

Answer 37

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

Question 38

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

Answer 38

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

Question 39

Explain the Aufbau principle.

Answer 39

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.

Question 40

Explain Pauli’s exclusion principle.

Answer 40

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.

Question 41

Explain Hund’s Rule.

Answer 41

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.

Question 42

Explain Stability of atomic orbitals.

Answer 42

(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.

Question 43

What is the formula for electronic configuration.

Answer 43

1s^2

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

Question 44

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

Answer 44

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.

Question 45

Define ionisation energy.

Answer 45

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.

Question 46

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

Answer 46

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

Question 47

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

Answer 47

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

Question 48

Define first ionisation energy.

Answer 48

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-

Question 49

Define the second ionisation energy.

Answer 49

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.

Question 50

Is Ionisation energy endothermic or exothermic?

Answer 50

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

Question 51

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

Answer 51

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.

Question 52

What are the factors which impact ionisation energy?

Answer 52

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

Question 53

What is the trend in ionisation energy across a period?

Answer 53

First Ionisation energy increases across a period.

Question 54

What is the trend in ionisation energy down a group?

Answer 54

First ionisation energy decreases down a group.

Question 55

Explain the trend in ionisation energy across a period?

Answer 55

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.

Question 56

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

Answer 56

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.

Question 57

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

Answer 57

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

Question 58

Explain the trend in ionisation energy down a group?

Answer 58

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.

Question 59

what is meant by successive ionisation energies?

Answer 59

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

Question 60

What is the trend in successive ionisation energies?

Answer 60

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.

Question 61

Define ground state.

Answer 61

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.

Question 62

Define excited state:

Answer 62

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

Question 63

Define isoelectronic.

Answer 63

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