Atomic Structure (electrostatic effects)

What are the main factors determining the strength of electrostatic attraction between the nucleus and electrons?

• Number of electron shells

• Nuclear charge

• Shielding effect

How does number of electron shells affect EFA between the nucleus and outermost electrons?

Number of electron shells increases → n of the outermost shell increases → number of inner electrons increases → EFA decreases

How does the number of protons affect EFA between the nucleus and outermost electrons?

Number of protons increases → Nuclear charge increases → EFA increases

How does shielding effect come about?

Contributed by inner shell electrons mainly as compared to outermost electrons

Define shielding

Repulsion between between electrons in the inner and outer shells preventing outer shell electrons from experiencing the full effect of the nuclear charge

Define the atomic radius

Half the shortest inter-nuclear distance found in the structure of the element

What is the trend of atomic radii across a period?

Nuclear charge increases as proton number increases → EFA increases → atomic radii decreases

Shielding effect remains approximately constant as electrons are added to the same outermost shell

What is the trend of atomic radii down a group?

• Number of electron shells increases

• Nuclear charge is less dominant as compared to other factors (expected trend)

• Shielding experienced by valence electrons increases significantly

EFA between nucleus and outermost electrons decreases → atomic radii increase

What is the trend of atomic radii across a period?

• Number of electron shells remains the same

• Nuclear charge increases

• Shielding is approximately constant as electrons are added to the same outermost shell

EFA between nucleus and outermost electrons increases → atomic radii decrease

Compare the radius of a cation to that of the parent atom

• Cation has 1 less electron shell

• Same nuclear charge

• Cation has fewer inner shell electrons

EFA between nucleus and outermost electrons for cations are stronger → smaller atomic radii

Compare the radius of an anion to that of the parent atom?

• Same number of electron shells

• Same nuclear charge

• Anion has more outer shell electrons

EFA between nucleus and outermost electrons for anions are weaker → larger atomic radii

What is the trend pf isolelectronic ions across a period?

• Same number of electron shells

• Nuclear charge increases across the period

• Same shielding with same number of electrons in cations and anions respectively

EFA between nucleus and outermost electrons increases → ionic radii decreases

Describe the sharp increase in ionic radius from Al3+ to P3-

• P3- has 1 more electron shell

• Nuclear charge is less dominant as compared to other factors (expected trend)

• P3- has more inner shell electrons than Al3+

EFA between nucleus and outermost electrons decreases sharply in P3-

Define 1st iionisation energy of M

The 1st ionisation energy is the energy required to remove 1 mole of gaseous M atoms to form 1 mole of gaseous M+ ions

2nd IE → M+ to M2+

Ionisation is endothermic

What are the equations for 1st and 2nd ionisation energies for M?

M (g) → Mg+ (g) + e-

M+ (g) → Mg2+ (g) + e-

Describe the trend in 1st IE across a period

• Number of electron shells remain the same

• Nuclear charge increases

• Shielding remains approximately constant

EFA between nucleus and outermost electrons increases → 1st IE generally increases

What are the 1st IE anomalies in G2 & 13

IE of Be > IE of B

IE of Mg > IE of Al

Explain the 1st IE anomalies in G2 & 13

• The p orbital electron to be removed is at a higher energy level than the s electron

• Less energy is required to remove the p electron as it is less attracted to the nucleus than the s electron

What are the 1st IE anomalies in G15 & 16?

IE of N > IE of O

IE of P > IE of S

Explain the 1st IE anomalies in G15 & 16?

S and P as examples:

• 3p electron removed from S is paired while that to be removed from P is unpaired

• Greater electron-electron repulsion between paired electrons → lower first IE

Describe the trend in 1st IE down a group

• Number of electron shells increases

• Nuclear charge is less dominant as compared to other factors (expected trend)

• Shielding increases with increasing number of inner shell electrons

EFA between nucleus and outermost electrons decreases → 1st IE generally decreases

Why is there a greater decrease in 1st IE between Ne and Na and between Ar and K?

• G1 element has 1 more electron shell than the G18 element

• Shielding experienced by valence electrons is greater in the G1 element

• Despite increasing nuclear charge, EFA decreases

Describe the trend of successive ionization energies?

• Same number of electron shells

• Same nuclear charge

• Shielding effect decreases as number of electrons decreases

EFA increases → increasing trend in successive IEs

How can group number be deduced from successive ionization data?

• Largest jump in IE

• Electron located in the inner electron shell that experiences less shielding and hence stronger EFA

Define electronegativity

Relative measure of an atom’s ability to attract bonding electrons

Describe the trend of electronegativity across a period

• Same number of electron shells

• Nuclear charge increases

• Shielding remains approximately constant

EFA increases → electronegativity increases

Describe the trend of electronegativity down a group

• Number of electron shells increases

• Nuclear charge is less dominant as compared to other factors (expected trend)

• Shielding increases with more inner electrons

EFA decreases → electronegativity decreases