D. Periodic Relationship Among Elements

Question 1

Explain effective nuclear charge (Zeff) and shielding effect

Answer 1

Nucleus (+) attracts electrons (-)
→ inner electrons experience stronger attractive force (closer to nucleus)
→ valence electrons experience weaker attractive force (due to shielding effect from inner electrons; farther from nucleus)

Question 2

Discuss the periodic trend of effective nuclear charge

Answer 2

1. Down a group: almost constant
   → although increases the nuclear charge, it also increases the shielding effect
2. Across a period: increases from left to right
   → increases the protons in nucleus
   → increases positive charge
   → increases nuclear attraction on outer electrons

Question 3

How to quantitatively determine the Zeff?

Answer 3

Zeff = Z - S

Z - nuclear charge/ Atomic Number
S - shielding constant

Question 4

A high Zeff indicates ____

Answer 4

strong attraction between nucleus and electron of interest

Question 5

What is Slater’s rule?

Answer 5

set of rules used to determine the shielding constant (S)
→ electrons that are closer to the nucleus contribute more significantly that those farther away

Question 6

Based on the first step of Slater’s rule, how will the electrons be grouped?

Answer 6

(1s) (2s, 2p) (3s, 3p) (3d) (4s, 4p) (4d) (4f) (5s, 5p) (5d) (5f) …

Question 7

Identify the shielding coefficient:

s and p
n → ___
n-1 → ____
n-2 → ____

d and f
n → ___
n-1 → ____
n-2 → ____

Answer 7

s and p
n → 0.35
n-1 → 0.85
n-2 → 1.00

d and f
n → 0.35
n-1 → 1.00
n-2 → 1.00

Question 8

In the problem, calculate the effective nuclear charge of a 5s electron in a Sr atom. Why only one electron from the 5s shell is included?

Answer 8

Electrons in the same shell do not shield each other. Same with electrons in higher energy level

Question 9

Differentiate three types of atomic radius

Answer 9

1. Ionic radius - half the distance between two nuclei of neighboring ions in an ionic compound
2. Metallic radius - half the distance between adjacent atoms in a metal lattice
3. Covalent radius - half the distance between nuclei in a molecule consisting of identical atoms

Question 10

Discuss the periodic trend of atomic radius

Answer 10

1. Down a group: atomic radius increases
   → additional energy levels (n increases)
   → increase distance between nucleus and outermost electrons
2. Across a period: atomic radius decreases
   → Zeff increases (proton increases)
   → greater attraction between nucleus and electrons
   → electrons are pulled

Question 11

Cations are ___ than their neutral atoms because ____

On the other hand, anions are ___ than their neutral atoms because _____

Answer 11

Cations = smaller
→ electrons are lost
→ reduces electron repulsion
→ allows nucleus to pull electron closer

Anions = larger
→ electrons are gained
→ increase in repulsion between added electron and inner electrons
→ electron cloud expands

Question 12

What is ionization energy?

Answer 12

Minimum energy required to remove an electron from a gaseous atom in its ground state

Question 13

Removal of electron using ionization energy is an ____ process and hence, ionization energy values are always ____

Answer 13

endothermic process
ionization energy → always positive

Question 14

Successive electron removal makes the atom _____ and hence, requires ____ ionization energies

Answer 14

Successive electron removal makes the atom MORE POSITIVE and hence, requires HIGHER ionization energies
→ inner electrons are held more tightly

Question 15

Discuss the periodic trend of ionization energy

Answer 15

1. Down a group: ionization energy decreases
   → distance between nucleus and outermost electrons increases
   → easier to remove farther, outer electrons
2. Across a period: ionization energy increases
   → Zeff increases
   → greater attraction between nucleus and electrons
   → harder to remove outer electron

Question 16

Define electron affinity

Answer 16

Energy change when an electron is added to a gaseous atom

Question 17

Negative electron affinity means ___

Positive electron affinity means ___

Zero electron affinity means ___

Answer 17

Negative electron affinity means energy is released when electron is added

Positive electron affinity means energy is gained when electron is added
→ less negative
→ lower affinity for electrons

Zero electron affinity means no change in energy when electron is added

Question 18

Discuss the periodic trend of electron affinity

Answer 18

1. Down a period: electron affinity decreases (more positive)
   → added electrons are farther from nucleus
   → weaker nuclear attraction
   → more energy is required
2. Across a period: electron affinity increases (more negative)
   → Zeff increases
   → stronger nuclear charge
   → stronger attractions to another electron
   → more energy is released

Question 19

Define electronegativity

Answer 19

ability of an atom to attract electrons; a relative measure that involves two atoms

Question 20

Discuss periodic trends of electronegativity

Answer 20

1. Down a group: electronegativity decreases
   → farther, outer electrons are weakly attracted to nuclear charge
2. Across a period: electronegativity increases
   → increase in no. of protons
   → Zeff pulls electrons more strongly

Question 21

Define metallic property

Answer 21

Tendency of an atom to lose electrons
→ lose electrons: metals
→ gain electrons: nonmetals

Question 22

Discuss periodic trends of metallic property

Answer 22

1. Down a group: metallic property increases
   → farther electrons, easier to lost
2. Across a period: metallic property decreases
   → ionization energy is higher
   → harder to lose electrons
   → higher electron affinity
   → easier to gain electron

Question 23

Summarize periodic trends

Answer 23

MA
↓←
(Metallic property, Atomic radius)

EEIZ
→↑
(Electron affinity, Electronegativity, Ionization energy, Zeff)

Question 24

Discuss the periodic trends of lanthanide

Answer 24

Lanthanide Contraction (exception to atomic radius trend)

→ atomic size decreases from left to right
→ poor shielding (spread out shape)
→ diffuse and big 4f inner electrons
→ increased Zeff, stronger attraction
→ applies to period 5 and 6

Question 25

Explain the exception of Al and Ga to the periodic trends

Answer 25

Ga has smaller atomic radius than Al → poor shielding of d-electrons from Ga → d-electrons are more complex → stronger attraction | also explains the irregular trend of transition metals

Question 26

Why atomic radius of periods 5 and 6 are similar?

Answer 26

due to lanthanide contraction

Question 27

Why noble gases have larger atomic radius?

Answer 27

Due to electron repulsion within fully filled outer shells

Question 28

Group 15 (nitrogen group) has higher _____ than group 16 (oxygen group) because ____

Answer 28

higher ionization energy of N than O → easier to remove electron from O due to repulsion of paired electrons (lower IE of O) → N is more stable because of weaker repulsion due to half-filled valence orbital (higher IE of N)

Question 29

Beryllium has ___ ionization energy than boron

Answer 29

higher ionization energy → harder to remove electron from stable, filled 2s valence orbital (higher IE of Be) → B has outermost, less tightly held electron in 2p orbitals (lower IE of B) | applies to whole group 13 (B group) and group 2 (Be group)

Question 30

Sodium has ____ electron affinity than magnesium

Answer 30

higher electron affinity (more negative, less positive) → release more energy when gaining electron (more negative) → added electron enters an s orbital (lower energy) | applies to whole group 1 and group 2

Question 31

Nonmetals in period 2 has ___ electron affinity than the nonmetals of period 3

Answer 31

lower electron affinity (less negative, more positive) → period 2 have stronger electron-electron repulsion due to smaller 2p orbitals → harder to accept electrons → period 3 has larger → accepts electrons more easily

Question 32

Arrange the atoms in increasing electron affinity Be, Mg, Ca, Sr

Answer 32

Be < Mg < Ca < Sr → as n increases, energy gap between filled ns2 and empty np orbital decreases → easier to accept electron into the np orbital

Question 33

Noble gases have ____ electronegativity than corresponding atoms in a period

Answer 33

lower or not electronegative at all → due to complete valence shell → only Kr and Xe