Atomic structure

Question 1

where do protons deflect to and where do electrons deflect to. describe the extent of deflection.

Answer 1

negative plate for protons and positive plates for electrons. electrons deflect to a larger extent compared to protons. angle of deflection is proportional to (charge)/(mass).

relative mass of a proton is 1 while the relative mass of an electron is

Question 2

define atomic number and nucleon number

Answer 2

atomic number refers to the no. of protons in an atom while nucleon number refers to the toal number of protons and neutrons in an atom.

atomic number is also** proton number**, while nucleon number is also **mass number **

Question 3

what is a nuclide

Answer 3

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Question 4

define isotopes.
define properties of isotopes

Answer 4

isotopes are atoms of the same element with the same number of protons but different number of neutrons.

isotopes possess similar chemical properties but different physcial properties because they have different relative abundances.

Question 5

define orbitals

Answer 5

they are regions of space around the nucleus where there is a 95% probability of locating the electron in this region

Question 6

there are 4 propertis

describes the properties of shells

Answer 6

1. each shell represents a discrete energy level in an atom
2. each shell is assigned a principal quantum number,n, which indicates the average distance of the orbitals from the nucleus.( n=1 means 1st shell closest to positively-charged nucleus)
3. higher the value of n, the higher the energy level of the electron
4. the maximum number of electrons that can be found in a given shell is 2n^2.

the last property is the most important

Question 7

properties of subshells

Answer 7

1. each shell is made up of at least one sub-energy level which are termed subshells.
2. the value of n will determine the number of subshells.
   3. generally, the order of the energy levels for the subshells is s<p<d<f.

Question 8

what are atomic orbitals

Answer 8

1. each subshell is a group of degenerate atomic orbitals(same energy) with different orientation in space.
2. each orbital can hold a maximum of 2 electrons.

Question 9

how many electrons can each subshell s,p,d,f hold?

Answer 9

S:2(1orb)
P:6(3orb)
d:10(5orb)
f:14(7orb)

Question 10

define s orbitals

Answer 10

they are non-directional and spherical in shape.
As n increases, shape of the S orbital remains unchaged with size of s orbital increases.

Question 11

define p orbitals

Answer 11

1. they are directional and dumb-bell in shape.
   2.there are different p orbitals(Px,Py,Pz). orbitals within a given subshell(2Px, 2Py,2Pz) are degenerate.
2. As n increases, shape of p orbitals remains unchanged but** orientation and size changes.**

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Question 12

define d orbitals

Answer 12

there are 5 different orbitals and they are all directional

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Question 13

energy level diagram(simplified)

Answer 13

no answer

an empty 4s subshell must be drawn lower than an empty 3d due to lower energy level

Question 14

What is the Aufbau(building-up principle)

Answer 14

electrons occupy the lowest enerfy orbital first before occupying the higher energy orbitals
empty 4s orbital filled up before 3d orbital.
However, if both are filled up, remove electrons from 4s before removing from 3d.

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Question 15

Hund’s rule

Answer 15

electrons are added into the orbitals of a subshell singly first with parallel spins before pairing takes place.

electrons possess two spin states, either spin up or spin down.

single electron is “unpaired” while two are electrons are “paired”.

for orbitals of a subshell which are degenerate, electrons are first positioned as far apart as possible to minimise inter-electronic repulsion.

Question 16

Pauli Exclusion Principle

Answer 16

each orbital can hold a maximum of two electrons in opposite spins.

Question 17

Anomalous electronic configurations” chromium and copper

Answer 17

chronium: [Ar] 3d^5, 4s^1
(more stable)

copper: [Ar] 3d^10, 4s^1
(more stable)

Question 18

what are isoelectronic species

Answer 18

they have the same number of electrons

Question 19

define first ionisation energy

Answer 19

energy required to remove one mole of electrons from one mole of gaseous atoms of the element to form one mole of singly positively charged gfaseous ions.

**ionisation energy valuesa re positive in sign as energy is absorbed to remove electron from a gaseous atom/ion. **

Question 20

define second ionsation energy

Answer 20

energy required to remove one mole of electrons from one mole of singly positively charged gaseous ions to form one mole of doubly positively charged gaseous ions.

max number of ionisation energies that an element can have is equal to its atomic no.

Question 21

what is nuclear charge

Answer 21

the electrostatic attraction between the protons in the nucleus of an atom and the surrounding electrons

as the no. of protons increases, nuclear charge increases.

attractive forces decreases for electrons that are found in orbitals further away from the atomic nucleus.

Question 22

what is screening effect

Answer 22

partial decrease in electrostatic attraction between the nucleus and its valence electrons due to repulsive forces from other electrons present

As no. of inner shells increases, screening effect increases.

there is minimal shielding between electrons from the same shell.

Question 23

what is effective nuclear charge

Answer 23

the net electrostatic attraction of the protons in the nucleus with the valence/outermost electrons.

effective nuclear charge = nuclear charge - screening effect

Question 24

explain why there is a general increase in successive ionisation energy values

Answer 24

1. since the no. of protons is constant, nuclear charge remains unchanged.

2.as electrons are being removed, the increasingly positive ion attracts its remaining electrons more strongly.

3. increasing amount of energy is required to remove each remaining valence electron, resulting in higher ionisation energy successively.

Question 25

explain why there is a significant increase in value from the first to second IE.

Answer 25

1. 2nd IE involves the **removal of an inner electron from an inner shell that experiences a stronger electrostatic attraction** with the nucleus compared to the outermost electron for the 1st IE. 2. **More energy **is required to remove the inner electron than the outermost electron, resulting in a higher value for 2nd IE. ## Footnote state which eletron belongs to which shell and orbital.

Question 26

why does atomic radius increase down the group

Answer 26

1. The **no. of shells increase, leading to an increase in screening effect**, where each successive element has its valence electrons located in a **shell with a higher principal quantum no., n**. 2. hence, valence electrons are** increasingly further away** from the nucleus, despite the **increase in nuclear charge**.

Question 27

why does first ionisation energy decrease down the group

Answer 27

1. The **no. of shells increase, leading to an increase in screening effect**, where each successive element has its valence electrons located in a **shell with a higher principal quantum no., n**. 2. hence, valence electrons are** increasingly further away** from the nucleus, despite the **increase in nuclear charge**. 3. a** smaller** amount of energy is required to remove the valence electron.

Question 28

why does atomic radius decrease across a period

Answer 28

1. across the same period, electrons are **added to the same valence shell **across the period, hence **screening effect remains approx. constant**. 2. **Increase in nuclear charge**, due to increase in number of protons,** outweighs negligible increase** in screening effect. 3. **effective nuclear charge increases** across the period, leading to **stronger electrostatic forces of attraction between nucleus and valence electrons**. 4. As a result, valence electrons are **progressively pulled closer** to the nucleus. ## Footnote however, argon does not follow the general decrease in atomic radius across period 3. this is due to the differences in measurement of atomic radius of an atom.

Question 29

why does first IE generally increase across the perios.

Answer 29

1. across the same period, electrons are **added to the same valence shell **across the period, hence **screening effect remains approx. constant**. 2. **Increase in nuclear charge**, due to increase in number of protons,** outweighs negligible increase** in screening effect. 3. **effective nuclear charge increases** across the period, leading to **stronger electrostatic forces of attraction between nucleus and valence electrons**. 4. **More** energy is required to remove the valence electron.

Question 30

EXCEPTIONS: why is the first IE of Mg greater than Al.

Answer 30

a smaller amont of energy is required to remove the 3p electron in Al which is of higher energy than the 3s electron in Mg.

Question 31

EXCEPTIONS: why is the first IE of P greater than S

Answer 31

1. there exists **inter-electronic repulsion** between the paired 3p electrons in S. 2. A **smaller **amount of energy is required to remove the paired 3p electron in S than the unpaired 3p electron in P. ## Footnote other things to note: 1. drawing of orbitals 2. the 3-8-8 method for graph analysis