Chapter 5 Periodicity and Atomic Structure. Development of the Periodic Table The periodic table is...

Chapter 5

Periodicity and Atomic Structure

Development of the Periodic Table

• The periodic table is the most important

organizing principle in chemistry.– Periodic table powerpoint – elements of a group have

similar properties– Chapter 2 – elements in a group form similar formulas

– Predict the properties of an element by knowing the properties of other elements in the group

Light and the Electromagnetic Spectrum

• Radiation (light) composed waves of energy

• Waves were continuous and spanned the electromagnetic spectrum

Light and the Electromagnetic Spectrum

Light and the Electromagnetic Spectrum

Light and the Electromagnetic Spectrum

• Speed of a wave is the wavelength (in meters)

multiplied by its frequency in reciprocal seconds.

Wavelength x Frequency = Speed

(m) x (s–1) = c (m/s–1)» C – speed of light - 2.9979 x 108 m/s–1

Electromagnetic Radiation and Atomic Spectra

• Classical Physics does not explain– Black-body radiation– Photoelectric effect– Atomic Line Spectra

• Blackbody radiation is the visible glow that solid

objects emit when heated.

• Max Planck (1858–1947): Developed a formula to

fit the observations. He proposed that energy is

only emitted in discrete packets called quanta.

• The amount of energy depends on the frequency:

E h

hc h 6.626 10 34 J s

Particlelike Properties of Electromagnetic Radiation: The Plank Equation

Particlelike Properties of Electromagnetic Radiation: The Plank Equation

• A photon’s energy must exceed a minimum threshold for electrons to be ejected.

• Energy of a photon depends only on the frequency.

Electromagnetic Radiation and Atomic Spectra

• Atomic spectra: Result from excited

atoms emitting light.

– Line spectra: Result from electron

transitions between specific energy

levels.

Electromagnetic Radiation and Atomic Spectra

1/λ = R [1/m2 – 1/n2]

Quantum Mechanics and the Heisenburg Uncertainty Principle

• Niels Bohr (1885–1962): Described atom as electrons

circling around a nucleus and concluded that electrons

have specific energy levels.

• Erwin Schrödinger (1887–1961): Proposed quantum

mechanical model of atom, which focuses on wavelike

properties of electrons.

Quantum Mechanics and the Heisenburg Uncertainty Principle

• Werner Heisenberg (1901–1976): Showed that

it is impossible to know (or measure) precisely

both the position and velocity (or the

momentum) at the same time.

• The simple act of “seeing” an electron would

change its energy and therefore its position.

Wave Functions and Quantum Mechanics

• Erwin Schrödinger (1887–1961): Developed a

compromise which calculates both the energy of

an electron and the probability of finding an

electron at any point in the molecule.

• This is accomplished by solving the Schrödinger

equation, resulting in the wave function, .

Wave Functions and Quantum Mechanics

• Wave functions describe the behavior of electrons.

• Each wave function contains three variables called

quantum numbers:

– • Principal Quantum Number (n)

– • Angular-Momentum Quantum Number (l)

– • Magnetic Quantum Number (ml)

Wave Functions and Quantum Mechanics

• Principal Quantum Number (n): Defines the size and

energy level of the orbital. n = 1, 2, 3,

• As n increases, the electrons get farther from the

nucleus.

• As n increases, the electrons’ energy increases.

• Each value of n is generally called a shell.

Wave Functions and Quantum Mechanics

• Angular-Momentum Quantum Number (l): Defines the three-dimensional shape of the orbital.

• For an orbital of principal quantum number n, the value of l can have an integer value from 0 to n – 1.

• This gives the subshell notation:

l = 0 = s orbital l = 1 = p orbital

l = 2 = d orbital l = 3 = f orbital

l = 4 = g orbital

Wave Functions and Quantum Mechanics

• Magnetic Quantum Number (ml): Defines the spatial orientation of the orbital.

• For orbital of angular-momentum quantum number, l, the value of ml has integer values from –l to +l.

• This gives a spatial orientation of:

l = 0 giving ml = 0

l = 1 giving ml = –1, 0, +1

l = 2 giving ml = –2, –1, 0, 1, 2, and so on…...

Wave Functions and Quantum Mechanics

Problem

• Why can’t an electron have the following quantum numbers?

– (a) n = 2, l = 2, ml = 1 (b) n = 3, l = 0, ml = 3

– (c) n = 5, l = –2, ml = 1

• Give orbital notations for electrons with the following quantum numbers:

– (a) n = 2, l = 1, ml = 1 (b) n = 4, l = 3, ml = –2

– (c) n = 3, l = 2, ml = –1

The Shapes of Orbitals

• s Orbital Shapes:

The Shapes of Orbitals

• p Orbital Shapes:

The Shape of Orbitals

• d and f Orbital Shapes:

Orbital Energy Levels in Multielectron Atoms

Orbital Energy Levels in Multielectron Atoms

• Zeff is lower than actual nuclear charge.

• Zeff increases toward nucleus ns > np > nd > nf

• This explains certain periodic changes observed.

Orbital Energy Levels in Multielectron Atoms

• Electron shielding leads to energy differences among orbitals within a shell.

• Net nuclear charge felt by an electron is called the effective nuclear charge (Zeff).

Wave Functions and Quantum Mechanics

• Spin Quantum

Number:

• The Pauli Exclusion

Principle states that no

two electrons can have

the same four quantum

numbers.x

Electron Configurations of Multielectron Atoms

• Pauli Exclusion Principle: No two electrons in

an atom can have the same quantum numbers

(n, l, ml, ms).

• Hund’s Rule: When filling orbitals in the same

subshell, maximize the number of parallel spins.

Electron Configurations of Multielectron Atoms

• Rules of Aufbau Principle:

1. Lower n orbitals fill first.

2. Each orbital holds

two electrons; each

with different ms.

3. Half-fill degenerate

orbitals before pairing

electrons.

Electron Configurations and Multielectron Atoms

Li 1s2 2s1

1s 2s

Be 1s2 2s2

1s 2s

B 1s2 2s2 2p1

1s 2s 2px 2py 2pz

C 1s2 2s2 2p2

1s 2s 2px 2py 2pz

Electron Configurations and Multielectron Atoms

N 1s2 2s2 2p3

1s 2s 2px 2py 2pz

O 1s2 2s2 2p4

1s 2s 2px 2py 2pz

Ne 1s2 2s2 2p5

1s 2s 2px 2py 2pz

S [Ne] [Ne] 3s2 3p4

3s 3px 3py 3pz

Problems

• Give the ground-state electron configurations for:

– Ne (Z = 10) Mn (Z = 25) Zn (Z = 30)

– Eu (Z = 63) W (Z = 74) Lw (Z = 103)

• Identify elements with ground-state configurations:

– 1s2 2s2 2p4 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2 4d6

– 1s2 2s2 2p6 [Ar] 4s2 3d1 [Xe] 6s2 4f14 5d10 6p5

Electron Configurations and the Periodic Table

Some Anomalous Electron Configurations

• Anomalous Electron Configurations: Result from unusual stability of half-filled & full-filled subshells.

• Chromium should be [Ar] 4s2 3d4, but is [Ar] 4s1 3d5

• Copper should be [Ar] 4s2 3d9, but is [Ar] 4s1 3d10

• In the second transition series this is even more pronounced,

with Nb, Mo, Ru, Rh, Pd, and Ag having anomalous

configurations (Figure 5.20).

Electron Configurations and Periodic Properties: Atomic Radii

Optional Homework

• Text – 5.24, 5.26, 5.28, 5.30, 5.32, 5.34, 5.44, 5.56, 5.58, 5.66, 5.68, 5.70, 5.72, 5.76, 5.78, 5.82, 5.84, 5.94, 5.98, 5.108

• Chapter 5 Homework online

Required Homework

• Assignment 5