atomic structure

Question 1

what is an atom?

Answer 1

an atom is the smallest component of an element having the chemical properties of the element.
all atoms are made up of electron, protons and neutrons.

protons & neutrons may be collectively referred to as nucleons

Question 2

what are isotopes?

Answer 2

isotopes are atoms of an element that have the same proton number but different number of neutrons (& electrons)

isotopes have same chemical properties (same no. of e) but different physical properties (diff no. of n)

Question 3

are protons deflected to the negative plate?

Answer 3

yes, protons (& cations) are deflected to the negative plate.
electrons (& anions) are deflected to the positive plate.
neutrons pass straight through

☆ if a particle is deflected to a smaller extent (smaller angle), it is because it is heavier as it has more particles. eg electrons have a much lower mass than protons so they are deflected to a greater extent. He has twice the electrons as H so He is deflected half as much as H

Question 4

what is relative atomic mass?

Answer 4

average mass of one atom of an element ÷ (1/12 x mass of one atom of ¹²C)

Question 5

what is relative isotopic mass?

Answer 5

mass of one atom of an isotope of an element ÷ (1/12 the mass of one atom of ¹²C)

Question 6

what is principal quantum number?

Answer 6

principal quantum number, n, describes the main energy level of an electron. larger n = higher energy level = electron is further away from nucleus.

maximum number of electrons that can occupy each principal quantum shell is given by 2n²

Question 7

what are the 3 rules for distributing electrons in various orbitals in an atom?

read through

Answer 7

1 - aufbau/’building up’ principle -> electrons occupy orbitals in order of energy levels; the orbital with the lowest energy is filled first (1s < 2s < 2p < 3s < 3p< 4s < 3f < 4p)
2 - hund’s rule of multiplicity -> each orbital must be singly occupied before electrons are paired
3 - an orbital cannot hold more than two electrons, and the 2 electrons sharing the same orbital must have opposite spins

group 1-2 -> s block
group 3-12 -> d block
group 13-18 -> p block (except helium which is 1s²)

Question 8

why does atomic radius increase down the group?

Answer 8

• down the group, no. of quantum shells increases → outermost electrons are further away from the nucleus → atomic radius increases
• nuclear charge & shielding effect increases down the group, so effective nuclear charge differs little down the group.

no. of quantum shells is the most important factor

Question 9

why does atomic radius decrease across a period?

Answer 9

• across a period, number of protons increases → nuclear charge increases
• shielding effect remains relatively constant as electrons are added to the same outermost shell
• (nuclear charge increases while shielding effect remains relatively constant so) effective nuclear charge increases → stronger electrostatic forces of attraction between the nucleus and outermost electron
• outermost electrons are pulled closer to the nucleus → decrease in atomic radius

Question 10

why does atomic radius remain relatively invariant across the first row of transition elements?

Answer 10

• no. of protons increases → nuclear charge increases
• electrons are added to inner 3d subshell which contributes to the shielding effect
• increasing shielding effect nullifies, to a considerable extent, the influence of each additional proton in the nucleus → effective nuclear charge remains almost constant → atomic radius is relatively invariant

Question 11

across the same period, why does ionic radius of cations decrease, then increase from last cation to first anion, then decrease among anions?

(ionic radius across same period) (eg period3)

Answer 11

• cations of Na, Mg, Al & Si are isoelectronic as they have the same no. of electrions (10). anions of P, S & Cl are isoelectronic as they all have 18 electrons.
• across the 2 isoelectronic series, nuclear charge increases while shielding effect remains the same (as no. of electrons remains the same)
• leads to increased effective nuclear charge and stronger attraction between the outermost electrons and nucleus → decreased ionic size across each series
• there is a sharp increase in ionic radius from cationic series to anionic series as anions have one more quantum shell than cations.

Question 12

what is first ionisation energy?

Answer 12

first ionisation energy is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of unipositively charged atoms

Question 13

what is the nth ionisation energy?

where n = 2, 3, 4, 5…

Answer 13

the nth ionisation energy of an element is the energy needed to remove 1 mole of electrons from 1 mole of gaseous ions with (n-1)+ charge to form 1 mole of gaseous ions wiith n+ charge

eg if there is a big increase from the 2nd to 3rd IE, it implies that the 3rd electron is in an inner quantum shell, so there are 2 electrons in the outermost or valence subshell (the atoms is grp2)

Question 14

why does first IE decrease down a group?

Answer 14

• down the group, no. of quantum shells increases → outermost electrons are further from the nucleus → weaker electrostatic forces of attraction between the nucleus and outermost electron → less energy is required to remove this electron
• both nuclear charge & shielding effect increases down the group → effective nuclear charge differs little down the group (so no. of quantum shells is the more important factor)

Question 15

why does first IE increase across period 2 & 3?

Answer 15

• across the period, nuclear charge increases while shielding effect remains relatively the same → effective nuclear charge increases → electrostatic forces of attraction between outermost electrons & nucleus become stronger → more energy is required to remove the outermost electron → first IE increases

Question 16

what are the 2 anomalies in trend of first IE across period 2 & 3 and why do they exist?

Answer 16

anomaly 1: small dip between Mg [1s² 2s² 2p⁶ 3s²] (grp 2) & Al [1s² 2s² 2p⁶ 3s² 3p¹] (grp 13). the 3p subshell is further away from the nucleus than the 3s subshell → weaker attraction between the nucleus & the outermost electron → less energy is required to remove the 3p electron from Al, resulting in a lower ionisation energy for Al compared to Mg

anomaly 2: small dip between P [1s² 2s² 2p⁶ 3s² 3p³] (grp 15) and S [1s² 2s² 2p⁶ 3s² 3p⁴] (grp 16). all the 3p electrons in P are unpaired. in S, 2 of the 3p electrons are paired → there is inter-electronic repulsion between the paired electrons in the 3p subshell in S → less energy than expected is required to remove one of these paired electrons from S

Question 17

why does first IE remain relatively invariant across transition elements?

Answer 17

• first IE involves removal of a 4s electron
• no. of protons increases → nuclear charge increases.
• electrons are added to the 3d subshell → shielding effect increases.
• shielding effect increases, thus nullifying, to a certain extent, the influence of each additional proton in the nucleus → effective nuclear charge remains almost constant → energy required to remove the outermost electron of each succeeding element remains relatively invariant.