Broderick's Chapter 2 atomic structure - atomic orbitals

Question 1

Degenerate Orbitals

Answer 1

atomic orbitals within a sub-level are of equal energy

Question 2

Why do electrons fill 4s before 3d ?

Answer 2

The basic principle electron occupies orbital with the lowest energy and 4s orbital have a lower energy than 3d orbitals.

Question 3

How are orbitals defined and how many electrons does each orbital have ?

Answer 3

a region of space where there is a high probability of finding an electron and each orbital contain maximum 2 electrons

Question 4

What is the Pauli exclusion Principle ?

Answer 4

There are two electrons in each orbital so therefore they must have opposite spins because they repel each other

Question 5

What is the Hund’s rule ?

Answer 5

every orbital in a subshell is singly occupied with one electron before it gets paired up

Question 6

What are the two exceptions of filling in orbital rules and explain them ?

Answer 6

The two exceptions are Chromium and Copper.
Copper and Chromium are exceptions to the common electron configuration methods because they are one electron away from reaching a more stable state (a half-filled 3d subshell for Chromium with 5 electrons under Hund’s Rule, and a filled 3d subshell with 10 electrons under the Pauli Exclusion Principle)

https://ask.learncbse.in/uploads/db3785/original/2X/7/7113372b4ca50ce5ec8be052c7aea8a0f18bd67b.png

Question 7

How to work out the condensed electronic configuration ?

Answer 7

Find the closest noble gas, group 0, and then count from that point onwards

e.g for Br use Br = [Ar] 3d^10, 4s^2, 4p^5

Question 8

First Ionisation energy

Answer 8

The energy required to remove one mole of electron from one of gaseous atoms under standard conditions

Question 9

What state are the elements in for ionisation energy ?

Answer 9

They have to be in gaseous state and state symbol has to be used

Question 10

Why is the second ionisation energy higher than the first ?

Answer 10

This is because the first removal of electron from a positive ion,

less electrons in the ion, making electron more tighter and closer to nucleus, stronger attraction

which requires more energy to break the attraction

Question 11

How do you interpret a Successive ionisation graph ? What do the big jumps tell you ?

Answer 11

It tells you which energy level or the electron that are being removed

Changes in subshell causes gradient to change

The big jumps tells us that the electron is removed from a shell closer to the nucleus

Question 12

Electrons in the ______ _____ do not _____ each other very well

Answer 12

same shell, shield

Question 13

How to answer ionisation graph questions ?

Answer 13

ALWAYS STATE Charge and distance

Question 14

What is the general trend for ionisation energy across a period ?

Answer 14

The ionisation energy increases from left to right across a period due to the increase in nuclear charge without a significant increase in shielding, greater force on outer electron

Question 15

Why is it easier to remove an electron from an orbital with two electrons than an orbital with one electron ?

Answer 15

This is because when there are two electrons in an orbital, there will be greater repulsion on the electrons thereby making it easier for it remove the electrons.

Question 16

Explain the trend of ionisation energy down a group

Answer 16

Ionisation energy decreases down a group

Despite nuclear charge increasing, the shielding also increases so the effective nuclear charge would be relatively the same

also the outer most electron is further away from the nucleus so lower electrostatic attraction between electron and nucleus = so less energy required to break this attraction

Question 17

How are electrons removed from transition metals ?

Answer 17

The 4s electrons are always removed before the 3d electrons

Question 18

What is the aufbau principle

Answer 18

electrons are fills subshell with lowest energy first

Question 19

Explain the general increase in successive ionization energies for the element.

Answer 19

The increase in successive IE is due to the increase in effective nuclear charge with the loss of each electron; there is a stronger attraction for remaining electrons for the same number of protons