Periodic Variations

Periodic trends

Predictable changes in element properties across the periodic table

Atomic size, ionization energy, and electron affinity

Main periodic properties discussed here are what?

They have similar valence-electron configurations

Why elements in the same group behave similarly

Increases down a group

Metallic character trend down a group

Covalent radius

Half the distance between nuclei of two identical bonded atoms

Increases

Atomic radius trend down a group

Decreases from left to right

Atomic radius trend across a period

Higher principal quantum number n; outer electrons are farther from nucleus

Why atomic radius increases down a group

Increasing effective nuclear charge pulls electrons inward

Why atomic radius decreases across a period

Picometers (pm)

Unit commonly used for atomic radius

64 pm

Fluorine covalent radius

99 pm

Chlorine covalent radius

114 pm

Bromine covalent radius

133 pm

Iodine covalent radius

Effective nuclear charge (Z_eff)

Net positive pull felt by an electron

Z_eff = Z - Shielding

Z_eff formula

Shielding

Reduction of nuclear attraction caused by other electrons

Core electrons

Which electrons shield most effectively?

It reduces attraction between nucleus and valence electrons

Why shielding matters

Generally increases from left to right

Trend of Z_eff across a period

Electrons are pulled closer to nucleus

Effect of larger Z_eff

Ionic Radius

Size of an ion.

smaller

Cation size compared to parent atom is

Loss of electrons increases effective nuclear pull on remaining electrons

Why cations are smaller

larger

Anion size compared to parent atom is

Extra electron-electron repulsion expands electron cloud

Why anions are larger

ionic radius

Trend of ___________ down a group increases

positive charge

Higher ______________ on cation causes radius to decrease

118 pm

Aluminum atom radius

68 pm

Al^(3+) ionic radius

104 pm

Sulfur atom radius

170 pm

S^(2-) ionic radius

Isoelectronic species

Atoms or ions with the same electron configuration

Number of protons

What determines size in an isoelectronic series?

nuclear charge

Greater ______________ in isoelectronic series causes radius to decrease

First ionization energy (IE_1)

Energy required to remove the first electron from a gaseous atom
X(g) → X^+(g) + e^-

Second ionization energy (IE_2)

Energy required to remove second electron
X^+(g) → X^2+(g) + e^-

Endothermic

Ionization processes are usually

Yes, it is true

Is it true that ionization energy trend across a period generally increases left to right.

Generally decreases

Ionization energy trend down a group

Valence electrons are farther from nucleus

Why larger atoms have lower ionization energies

Electrons are held more tightly

Why smaller atoms have higher ionization energies

Boron loses a higher-energy p electron

Why boron has lower IE than beryllium

Removing paired electron reduces electron-electron repulsion

Why oxygen has lower IE than nitrogen

Half-filled p subshell is especially stable

Nitrogen stability reason

Successive ionization energies

Energies required to remove electrons one after another

always increase

General trend for successive ionization energies

Core electron removal begins

Large jump in ionization energy means

Positive ion attracts remaining electrons more strongly

Why removing electron from cation is harder

Electron Affinity (EA)

Energy change when an electron is added to a gaseous atom.

Negative electron affinity

means Energy released when electron is added.

Positive electron affinity

means Energy required to add electron

more negative

Electron affinity trend across a period generally becomes _____________ left to right

Adding electron requires entering higher-energy shell

Why noble gases have unusual EA values

Their ns subshell is already filled

Why group 2 elements show EA exceptions

Half-filled np subshell gives extra stability

Why group 15 elements show EA exceptions

Chlorine

Element with most negative electron affinity

−348 kJ/mol

Chlorine electron affinity