Atomic Structure And Ionisation Energy

Question 1

What is an atom

Answer 1

The smallest part of an element That can take part in a chemical reaction

Question 2

What are the roles of protons, neutrons and electrons

Answer 2

Proton- determines identity of the element

Neutron- provides stability

Electrons- responsible for reactivity

Question 3

Describe the structure of the atom

Check table of subatomic particles

Answer 3

Atoms are mostly empty space surrounding a very small, dense nucleus that contains protons and neutrons (together called nucleons). Electrons are in shells in the empty space surrounding the nucleus.

•The nucleus of an atom is surround by electrons that occupy shells, or orbitals of varying energy levels.

Note: the nucleus is positively charged and negatively charged electrons orbit the nucleus due to the attraction between them.

Question 4

Describe the effects of electric fields on subatomic particles

Answer 4

If a stream of subatomic particles is pilasses through an electric field, they are seperated since they have different charges.

1. The electrons deflect towards the positive plate
2. Protons deflect towards the negative plate
3. Neutrons are undeflected

Note: the degree of deflection of electrons is more than That of protons since electrons are lower on mass.

Question 5

What are subshells (subsidiary quantum shells)?

Answer 5

Region of the principal quantum shells where electron exist in defined areas associated with particular amounts of energy.

S,p,d,f..

A collector of subshells of similar energy is called a SHELL or energy level

Question 6

What’s the relationship between the number of subshells into which a shell is divided?

Answer 6

Number of subshells= principal quantum number of the shell

Note: the farther we move from the nucleus, the greater the energy state of the shells/subshells

Question 7

What is an orbital? Check shape

Answer 7

A region of space outside the nucleus where an electron is most likely to be found (95% chance) at any moment; can be occupied by a MAX OF TWO ELECTRONS

NOTE: Shell> subshell>orbital

1. In the s subshell electrons occupy the s orbital
2. In the p subshell electrons occupy the p orbital etc

C

Question 8

Check Aufbaw law

Answer 8

Note: in general, while moving away from the nucleus the relative energy of the subshells increases.

1s, 2s 2p, 3s 3p 4s 3d, 4p

Question 9

Distribution of electrons in atoms

Answer 9

When electrons are placed in an atom they must first fill the lowest energy subshell (1s closest to nucleus). Afterwards, they fill the next available subshell stepwise until all the electrons are distributed.

Question 10

Why is the 4s subshell filled before the 3d subshell

Answer 10

The 4s orbital is filled before the the 3d subshell because electrons in the 4s subshell have a lower energy state

Note:
1.when an electron enters the 3d subshell electrons in the 4s subshell gain energy and they are placed after

2. Check summarized form:
   2He : 1s2
   10Ne: 1s2 2s2 2p6
   18 Ar: 1s2 2s2 2p6 3s2 3p6

Question 11

What are the steps when forming cations?

Answer 11

1. Electrons are always lost from the MOST ENERGETIC subshell (outer shell)
2. P electrons are removed first, then 2 electrons and then d electrons
3. Paired electrons are removed before unpaired electrons in the same sub energy level.

Question 12

What are the steps taken to determine the possible oxidation states or a transition element?

Answer 12

1. All electrons are first removed from the 4s subshell
2. Electrons are then removed stepwise from the 3d subshell until it is either:
   - half-filled or
   - empty

Question 13

Distribution of electrons in anions

Answer 13

.

Question 14

Check box and arrow diagrams and describe the discrepancies of copper and chromium

Answer 14

1. They have ONLY 1 electron in the 4s subshell
2. The 3d subshell is either half-filled (chromium) or fully-filled ( copper). This confers extra stability to their structures.

24Cr: (Ar) 3d5 4s1
29Cu: (Ar) 3d10 4s1

Question 15

What is a free radical and how are they formed?

Answer 15

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

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.

Question 16

Describe the position of elements in their groups in terms of ‘blocks’

Answer 16

Group 1 and 2: S block
Group 13 to 18: p block
Group 3 to 13: d block (elements that add electrons to the d subshell)

Question 17

Define the first ionisation energy

Answer 17

The energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous unipositive ions.

Question 18

Define successive ionisation energy

Answer 18

The energy required in each step to remove the first electron, then the second, then the third, and so on, from a gaseous atom.

Question 19

Define the second and third ionisation énergies

Answer 19

Second I.E: the energy required to remove one mole of electrons in one mole of gaseous unipositive ions to give one mole of gaseous bipositive ions.

Third I.E: the energy required to remove one mole of electrons from one mole of gaseous bipositive ions to form one mole of tripositove ions.

Question 20

What is the general formula for the Nth I.E?

Answer 20

M(n-1)+ (g) —> M(n)+ (g) + e-

The charge on the RHS gives the number of I.E

Question 21

What is the unit of ionisation energy?

Answer 21

kJ/mol

Question 22

What are the factors affecting the ionisation energy?

Answer 22

1. Size of nuclear charge
2. Distance of outer electrons from the nucleus
3. Shielding effect of inner electrons
4. Spin-pair repulsion

Question 23

How does the size of the nuclear charge affect the ionisation energy?

Answer 23

In general, I.E increases as the nuclear charge/proton number increases.

• As the atomic number increases, the positive nuclear charge in the nucleus increases
• The greater the positive nuclear charge, the greater the attractive force between the nucleus and the outer electrons.
• Hence, more energy is required to overcome such forces if an electron is to be removed.

Question 24

How does the distance of the outer electrons from the nucleus affect the I.E?

Answer 24

In general, the further an electron is from the nucleus, the lower the I.E

• The electrostatic force of attraction between the positive nucleus and negatively charged electrons decreases rapidly as the distance between them increases.
• Therefore, electrons further away from the nucleus are less attracted than those closer to the nucleus.

Question 25

What is shielding by inner electrons?

Answer 25

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

Question 26

How does shielding affect the I.E?

Answer 26

In general, the I.E is lowered as the number of full electron shells between the Nucleus and the outer electrons increases.

The greater the shielding effect if inner electrons on outer electrons, the lower the attractive force between the nucleus and the outer electrons.

Explanation:
• electrons tend to repel each other as they are all negatively charged.
• electrons in full inner shells repel electrons in the outer shells.
• therefore full inner shells of electrons reduce the refactor of the nuclear charge on the outer electrons; outer electrons only feel the residual nuclear charge)

Question 27

What is spin-pair repulsion?

Answer 27

A pair of electrons in the same orbital repel each other because they have the same charge.

Note:
1.electrons are placed in anti parallel directions within an orbital to reduce the repulsion between them
2. Spin-pair repulsion > repulsion of single electrons in separate orbitals.
This is why electrons are placed in separate orbitals before being paired up.

Question 28

How does spin-pair repulsion affect the I.E?

Answer 28

An increase in repulsion, makes it easier to remove an electron. Thus, the ionisation energy is lowered.

• Electrons in the same atomic orbital repel each other more than electrons in separate orbitals.

Question 29

Interpreting successive ionisation energies.

E.g Na

Answer 29

1. The first electron removed has a low first I.E: farthest from the nucleus and well shielded by inner electron shells.
2. Big jump in value of I.E between 1st and 2nd I.E: change in shells. Na has one valence electrons
3. From 2nd to 9th I.E there’s a GRADUAL increase in successive I.E: electrons are in the same shell.
4. 10th and 11th electrons have extremely high I.E:
   •in shell closest to the nucleus
   • no inner electrons for shielding
   • great force of attraction between nucleus and these electrons

Note:
1.a large increase in successive I.E suggest a change in shells.
2. Recall how to interpret group number and period number
• group number: number of valence electrons
• period number: number of electron shells

Question 30

Why is the variation of values of first I.E of elements considered a periodic property?

Answer 30

When plotted against atomic number, it shows a repeating pattern.

Question 31

Describe the trend in ionisation energy across a period

Answer 31

There is a gradual increase in ionisation energy across a period.

Explanation:
1.When moving across the period the size of the positive nuclear charge increases. 2. However the electrons are removed from the same shell. 3. The force of attraction between the positive nucleus and the outer electrons increases because:
• the size of the nuclear charge Increaes.
• The distance between the nucleus and the outer electrons remains reasonably constant.
• the shielding by inner electrons remains the same.

Question 32

Why is there a RAPID DECREASE in I.E between the LAST ELEMENT in one period and the FIRST ELEMENT in the next period.

Answer 32

1. When an electron is removed from the first element of the next period, it is being removed from a successively higher principal quantum shell, which is further away from the nucleus, compared to the last element in the last period. 2. The force of attraction between the positive nucleus and the outer electrons decreases because:
   • the distance between the nucleus and the outer electrons increases.
   • the shielding effect by inner electrons increases.
   • together these two factors outweigh the effect of the increase in nuclear charge.

Question 33

Why is their a discrepancy in this general trend between group 2 and group 13 ( e.g Be and B)

Answer 33

Be: 1s2 2s2
B: 1s2 2s2 2p1

The fifth electron in B is in the p subshell which is slightly further away from the nucleus than the 2s subshell.
There is less attraction between the fifth electron in B and the nucleus because:
• the distance between the Nucleus and the outer electron increases slightly
• the shielding effect by inner shell/subshell increases slightly.
• these two factors outweigh the increased nuclear charge.

Question 34

Explain the discrepancy in the I.E between group 15 and Group 16 elements (e.g N and O)

Answer 34

In N, an electron is removed form an unpaired orbital. Moreover, half-filled subshells confer extra stability to the atom. However, in O the electron is removed form an orbital that contains a pair of electrons. The extra repulsion between the pair of electrons in this orbital results in less energy needed to remove an electron, despite the higher atomic number of O.

Thus, the first I.E of O is lower because of spin-pair repulsion.

Question 35

Describe and explain the general trend in First I.E when moving down a group.

Answer 35

The first I.E decreases when moving down a group.

Explanation: when moving down a group, the electron is removed from the same type of orbital but from a successively higher principal quantum level. Although the nuclear charge increases down the group, there is less attraction between the positive nucleus and the outer electrons because:
•the distance between the nucleus and the outer electron increases.
•the shielding effect by complete inner shells increases.
•these two factors outweigh the increased nuclear charge.

Question 36

Define the covalent atomic radius.

Answer 36

The covalent atomic radius is half the distance between the nuclei of two covalently bonded atoms of the same type.

Question 37

Why are stable cations smaller than their corresponding atoms?

Answer 37

Atoms lose their outer shell electrons to form cations. Hence, the attractive force between the nucleus and the outer electron of the cation is larger. Electrons are pulled closer towards the nucleus.

Question 38

Why are stable anions larger than their corresponding atoms.

Answer 38

Atoms gain electrons to complete their outer shells, resulting in the formation of anions. Therefore, the attractive forces between the nucleus and outer electrons in anions is smaller. The outer electrons are further away from the nucleus.

Question 39

Describe and explain the trend in atomic radii down a group.

Answer 39

The atomic radius increases down any group.

Explanation:
When moving down the group, each successive element has one more shell of electrons which is further away from the nucleus. Although their is an increase in nuclear charge on going down the group, the increased shielding effect of inner shell electrons outweighs the increase in nuclear charge.
• weaker force of attraction between positive nucleus an outer electrons- the latter is further away.

Question 40

Describe and explain the trend in atomic radii across a period.

Answer 40

The atomic radius decreases across any period.

Explanation:

1. When moving across a period, the proton number and hence the nuclear charge increases regularly.
2. However the extra electrons are being added to the same shell.
3. Hence, the shielding effect by inner shell electrons does not change significantly.

This results in a greater attractive force between the the positive nucleus and the outer electrons which are pulled closer to the nucleus.

Question 41

Describe and explain the trend in Ionic radii down a group.

Answer 41

The ionic radius increases down any group.

1. In each group the number of electrons gained/lost by each element is generally the same.
2. Going down any group, each successive element has one more shell of electrons which is further away from the nucleus. The increased shielding effect of inner shell electrons outweigh the effect of the increased nuclear charge.
3. Outer electrons are less strongly attracted and hence further away from the nucleus.

Question 42

Describe and explain the trend in ionic size across a period from Group 1 to Group 14

Answer 42

Ionic radius decreases.

Explanation:
The increasing nuclear charge affects the outer electrons closer to the nucleus with increasing atomic number. The shielding effect of inner shell electrons reliait approximately the same ( lower effect than the increase in nuclear charge)

Note: this change on ionic radii is generally GREATER than that of the corresponding atomic radii due to the INCREASE IN CHARGE when moving from group 1 to group 14.

Question 43

Describe and a explain the trend in ionic radii across any period from group 15 to 18.

Answer 43

Ionic radii decreases (same reason as the decrease in atomic radii) + normal charge on the ions increases.

Question 44

Describe the relative charges of electrons and protons

Answer 44

The charge of a single electron is -1.602 x 10-19 coulombs, whereas the charge of a proton is +1.602 x 10-19 coulombs.
So, relative to each other, their charges are -1 and +1 respectively

Question 45

Describe mass number and atomic number.

Answer 45

•The ATOMIC NUMBER (or proton number) is 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
E.g. the atomic number of lithium is 3, meaning that a neutral lithium atom has 3 protons and therefore, also has 3 electrons

• The mass number (or nucleon number) is the total number of protons + neutrons in the nucleus of an atom, and has the symbol A
• The number of neutrons can be calculated by:

Number of neutrons = mass number – atomic number

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

Question 46

What is the principal quantum number?

Answer 46

•The principal quantum number indicates the energy level of a particular shell but also indicates the energy of the electrons in that shell.
A 2p electron is in the second shell and therefore has an energy corresponding to n = 2

Question 47

Why do electrons occupy the same region of space in orbitals despite their repulsion? (Spin-pair)

Answer 47

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