atomic structure

Question 1

what is the mass of an electron relative to the mass of a proton?

Answer 1

1/1836

Question 2

why does the electron have a greater extent of deflection?

Answer 2

this is because the electron has a greater charge-to-mass ratio

Question 3

how to calculate the angle of deflection

Answer 3

k (charge/mass)

Question 4

define mass number

Answer 4

it is the total number of protons and neutrons in the nucleus of an atom

Question 5

define atomic number

Answer 5

it is the number of protons in the nucleus of an atom

Question 6

define isotopes

Answer 6

they are atoms of the same element with the same number of protons and electrons but with a different number of neutrons.

Question 7

what are isoelectronic species?

Answer 7

it is a group of atoms and ions having the same number of electrons

Question 8

what is an alpha particle?

Answer 8

it is basically a helium atom

Question 9

what is a beta particle?

Answer 9

it is basically releasing an “electron”, causing the atomic number to increase by 1.

Question 10

what happens as the principal quantum number (n) increases?

Answer 10

the orbital becomes larger, the electron density is further away from the nucleus, the electrons have higher energy, and a weaker electrostatic attraction between the nucleus and the electron.

Question 11

what is an atomic orbital?

Answer 11

it is a region of space in which there is a 95% probability of locating the electron residing in it.

Question 11

what is a ‘p’ orbital?

Answer 11

it is dumb-bell in shape and directional along the x, y, z axes. for the same n, they are degenerate. as n increases, orbital size increases and the p orbital becomes more diffuse.

Question 12

what is an ‘s’ orbital?

Answer 12

it is spherical in shape and non-directional. as n increases, the orbital size increases and the ‘s’ orbital becomes more diffuse.

Question 12

what is a ‘d’ orbital?

Answer 12

it has 5 orbitals which are degenerate and 4 of them look like a clover leaf.

Question 13

state electronic configuration

Answer 13

1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p…

Question 14

what is hund’s rule?

Answer 14

electrons of parallel spins occupy singly in orbitals of a subshell in order to minimise inter-electronic repulsion.

Question 15

how does nuclear charge affect the strength of electrostatic attraction between the nucleus and electrons?

Answer 15

as the number of protons increases, the nuclear charge increases which means that electrostatic attraction increases.

Question 16

how does the shielding effect affect the strength of electrostatic attraction between the nucleus and electrons?

Answer 16

when the number of inner shell electrons increases, shielding effect experienced by the valence electrons increases which results in electrostatic attraction to decrease.

Question 17

how does the number of principal quantum shells affect the strength of electrostatic attraction between the nucleus and electrons?

Answer 17

the distance between the nucleus and the valence electrons increases which means that electrostatic attraction decreases.

Question 18

how to calculate effective nuclear charge?

Answer 18

subtract shielding effect from nuclear charge

Question 19

define first ionisation energy.

Answer 19

the first IE or an element is defined as the energy absorbed to remove one mole of electrons from one mole of gaseous atoms to form one mole of singly charged positive gaseous ions.

Question 20

what is atomic radius?

Answer 20

it is half the shortest inter-nuclear distance found in the structure of the element.

Question 21

define electronegativity

Answer 21

it is a measure of the ability of an atom in a molecule to attract a shared pair of electrons in a covalent bond.

Question 22

why does electrostatic attraction increase across a period?

Answer 22

across a period, there is an increasing nuclear charge due to an increase in the number of protons. there is an approx. constant shielding effect as additional electrons are added to the same valence shell. effective nuclear charge hence increases and the valence electrons experience stronger nuclear attraction.

Question 23

how does 1st IE change across a period?

Answer 23

since more energy is required to remove the valence electron due to stronger electrostatic attraction, first IE increases.

Question 24

how does atomic radius change across a period?

Answer 24

due to stronger electrostatic attraction, electrons are pulled closer to the nucleus of an atom. hence, atomic radius decreases.

Question 25

how does electronegativity change across a period?

Answer 25

with stronger nuclear attraction, the ability of an atom to attract a shared pair of electrons in a covalent bond increases which means electronegativity increases.

Question 26

why does electrostatic attraction decrease down the group?

Answer 26

down a group, there is an increasing nuclear charge due to an increase in the number of protons. there is an increasing shielding effect as well due to an increase in the number of inner shell electrons. thus, effective nuclear charge remains constant. valence electrons are further away from the nucleus due to an increase in the number of principal quantum shells, thus valence electrons experience weaker nuclear attraction.

Question 27

how does 1st IE change down the group?

Answer 27

with a weaker nuclear attraction, less energy is required to remove the valence electron. thus, 1st IE increases down the group.

Question 28

how does atomic radius change down the group?

Answer 28

with a weaker nuclear attraction and an increase in the number of principal quantum shells, atomic radius increases down the group.

Question 29

how does electronegativity change down the group?

Answer 29

with a weaker nuclear attraction, the ability of an atom to attract a shared pair of electrons in a covalent bond decreases which means that electronegativity decreases down the group.

Question 30

why is the 1st IE of Al lower than the 1st IE of Mg? (irregularity: between ns2 and ns2 np1)

Answer 30

this is because the electrons in the 3p-subshell are further away from the nucleus than the electrons in the 3s subshell and experience weaker nuclear attraction. thus less energy is required to remove the 3p electron in Al than the 3s electron in Mg.

Question 31

why is the 1st IE of oxygen lower than the 1st IE of nitrogen? (between ns2 np3 and ns2 np4)

Answer 31

inter-electronic repulsion between paired electrons in the same orbital exists so less energy is required to remove the 3p electron from oxygen than from nitrogen.

Question 32

how to determine the group of an element using ionisation energies?

Answer 32

find a large increase between ionisations energy.

Question 33

define ionic radius

Answer 33

it is the radius of a spherical ion in an ionic compound.