Chapter 3 Periodic Properties

Periodic Law

The principle that physical and chemical properties of elements repeat predictably when arranged by increasing atomic number.

Mendeleev's arrangement of elements

Elements were arranged by increasing atomic mass and similar properties.

Gaps in Mendeleev's table

Mendeleev left gaps to predict undiscovered elements would fill them.

Successful predictions by Mendeleev

He predicted elements such as gallium (Ga) and germanium (Ge).

Modern periodic table arrangement

Elements are arranged by increasing atomic number.

Pauli Exclusion Principle

No two electrons in the same atom can have identical sets of quantum numbers.

In 1s², what does 1 represent?

It represents the energy level.

In 1s², what does s represent?

It represents the type of orbital.

In 1s², what does ² represent?

It represents the number of electrons in those orbitals.

Orbital diagram box

Represents one orbital.

Orbital diagram arrow

Represents an electron.

Direction of arrow in orbital diagram

It represents the relative electron spin.

Hund’s Rule

Equal-energy orbitals fill singly first with parallel spins.

Which fills first: 4s or 3d?

4s fills first.

Why does 4s fill before 3d?

4s is lower in energy.

Valence electrons

Outermost electrons important in chemical bonding.

Core electrons

Electrons in complete inner energy levels.

Shielding/screening

The repulsion of an electron by other electrons from the full nuclear charge.

Effective nuclear charge formula

Zeff = Z − S.

Z in Zeff = Z − S

Z is the atomic number.

S in Zeff = Z − S

S is the screening constant, usually close to the number of inner electrons.

Greater penetration toward nucleus means

The electron experiences more nuclear attraction.

Why are there different energy sublevels?

Shielding and penetration affect electron energy.

Periodic table blocks correspond to

Different orbital types.

s-block

The left side of the periodic table.

p-block

The right side of the periodic table.

d-block

Transition metals.

f-block

Inner transition metals.

Condensed electron configuration

Uses the previous noble gas in brackets.

Se condensed configuration

[Ar] 4s² 3d¹⁰ 4p⁴.

Bi condensed configuration

[Xe] 6s² 5d¹⁰ 4f¹⁴ 6p³.

Cu electron configuration exception

[Ar] 4s¹ 3d¹⁰.

Why does Cu have an unusual configuration?

4s and 3d orbitals are very close in energy.

Metals location on periodic table

Lower left and middle of the periodic table.

Metal properties

Shiny, malleable, ductile, good conductors.

Metals tend to

Lose electrons.

Nonmetals location on periodic table

Upper right.

Nonmetal properties

Often gases or dull/brittle substances; poor conductors.

Nonmetals tend to

Gain electrons.

Metalloids location on periodic table

Along the stair-step line.

Metalloids properties

Properties of both metals and nonmetals.

Many metalloids are

Semiconductors.

Group 1A elements

Alkali metals.

Group 2A elements

Alkaline earth metals.

Group 7A elements

Halogens.

Group 8A elements

Noble gases.

Alkali metal valence configuration

ns¹.

Alkali metals usually form

+1 ions.

Alkaline earth metals have

ns².

Alkaline earth metals usually form

+2 ions.

Halogens need how many electrons for a full shell?

1 electron.

Halogens usually form

−1 ions.

Noble gases valence electrons

8, except He has 2.

Why are noble gases unreactive?

Their electron configurations have full, especially stable shells.

What largely determines an element's chemical properties?

Number of valence electrons.

Bonding atomic radius

Half the distance between two covalently bonded nuclei.

Atomic radius trend across a period

Decreases.

Atomic radius trend down a group

Increases.

Why does atomic radius decrease left to right?

Increasing Zeff pulls electrons closer.

Why does atomic radius increase down a group?

Increasing principal energy level n.

Where are the largest atoms located?

Bottom-left of the periodic table.

Cation compared with parent atom

Smaller.

Why is a cation smaller?

It loses valence electron(s).

Anion compared with parent atom

Larger.

Why is an anion larger?

It gains electron(s).

Ionic size trend down a group

Increases.

Ionization energy (IE) definition

Energy required to remove an electron from a gaseous atom or ion.

First ionization energy (IE₁)

Energy required to remove the first electron.

Successive ionization energies

Get progressively larger.

Huge jump in IE indicates

All valence electrons have been removed and core electrons are being removed.

Alkali metals have a huge jump at

IE₂.

Alkaline earth metals have a huge jump at

IE₃.

Ionization energy trend across a period

Generally increases.

Ionization energy trend down a group

Decreases.

Why does ionization energy increase left to right?

Zeff increases.

Why does ionization energy decrease going down?

Valence electrons are farther from the nucleus.

IE exception: Group IIA to IIIA

Electron is removed from a p orbital, which is farther from the nucleus.

IE exception: Group VA to VIA

Electron is removed from a doubly occupied orbital, where electron-electron repulsion makes removal easier.

Electron affinity (EA) definition

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

Example of electron affinity

Cl(g) + e⁻ → Cl⁻(g).

Electron affinity trend left to right

Generally becomes more exothermic.

EA exception: Group 1A to 2A

Added electron must enter a p orbital instead of an s orbital.

EA exception: Group 4A to 5A

Added electron must enter an already occupied orbital, causing repulsion.

Atomic Radius Trend

← Increases; ↓ Increases.

Ionization Energy Trend

→ Increases; ↑ Increases.

Metallic Character Trend

← Increases; ↓ Increases.

Electron Affinity Trend

Generally becomes more exothermic.

Memory Trick for Trends

Bottom left = Big + Metallic; Top right = Tight + Hard to remove electrons.