Chapter 8 Periodic Properties of the Elements. Energy of an atomic orbital For an atom, electrons...

Chapter 8

Periodic Properties of the Elements

H E

Energy of an atomic orbital

For an atom, electrons are in atomic orbitals.

Orbital Energy Levels for the Hydrogen Atom

H atom: E only depends on n degenerate

E depends on n and l

same n, l↑ ↔ E↑

H E

i fE E hv

A Picture of the Spinning Electron

Spin quantum number ms

ms = +1/2 or −1/2

4 quantum numbers are used to specify an electron.

How do electrons fill up atomic orbitals?

Pauli Exclusion Principle

In a given atom, no two electrons can have the sameset of four quantum numbers.

An orbital can hold only two electrons, they must haveopposite spins.

1s1

2s1

2p1

electron configuration

Lowest energy: ground state

Excited states

↑1s

↓2s

H atom

orbital diagram

↑2p

Now we can write the ground state electronconfigurations and draw orbital diagrams accordingto Pauli principle.

Electron configurations explain many chemicalproperties.

For degenerate orbitals, the lowest energy is attainedwhen the number of electrons with the same spin is maximized.

Hund’s rule

Valence electrons: electrons in the outermost shell formain group elements.They are involved in bonding and reactions.

Core electrons: inner electrons

Elements in the same group have similar valenceelectron configuration — similar chemical properties.

Noble gases have 8 (He 2) valence electrons. Stablestructure.

Number of valence electrons = main group number

Metals: tend to lose valence electrons to reach 8(2) valence electron.

Nonmetals: tend to gain electrons to reach 8(2) valence electrons.

Number of filled shells = period number

Review Problem Set 10

Periodic trends in atomic properties

• Atomic radius

Atomic Radii (in Picometers) for Selected Atoms

Atomic radius

In a period: decreases from left to right

In a group: increases from top to bottom

On the basis of periodic trends, choose the larger atom in each pair (if possible). Explain your choices.

(a) N or F (b) C or Ge (c) N or Al (d) Al or Ge

EXAMPLE 8.5 Atomic Size

Choose the larger atom or ion from each pair.

(a) S or S2– (b) Ca or Ca2+ (c) Br– or Kr

EXAMPLE 8.7 Ion Size

Periodic trends in atomic properties

• Atomic radius

• Ionization energy

Energy required to remove an electron froma gaseous atom or ion.

X(g) → X+(g) + e−

X+(g) → X2+(g) + e−

Ionization energy

first ionization energy

second ionization energy

Ionization energy

In a period: increases from left to right

In a group: decreases from top to bottom

(general trend)

Periodic trends in atomic properties

• Atomic radius

• Ionization energy

• Electron affinity

Energy change associated with the additionof an electron to a gaseous atom.

X(g) + e− → X−(g)

Electron affinity

∆E = Ef − Ei = EA < 0

X(g) + e−

X−(g)

EEi

Ef

stable X−

Electron affinity

In a period: increases from left to right

In a group: no clear trend

(very rough trend)

Periodic trends in atomic properties

• Atomic radius

• Ionization energy

• Electron affinity

Remember the trends

Problem Set 11