Chapter 3 The Structure of the Atom

In order to explain much of what is observed in chemistry, we need to adopt a model for the atom where the atom has components:

• Protons• Neutrons• Electrons

(We do not need to account for any structure within these components.)

Rutherford’s Model of the Atom

● Mass of atom is concentrated in nucleus.● Volume of atom is “empty.”● Model is still used today.

Particles and Waves

● Light as a waveWavelength, λ

Speed, c

Frequency, ν● c=λν

Particles and Waves

● Particles have mass and volume.● Waves have λ, ν and speed.

Light and Other Forms of Electromagnetic Radiation (EMR)

● Visible light is one form of EMR.● Other forms include:

Microwaves

X rays

Ultraviolet light● In all cases

c = νλ

Atomic Spectra

● Emission spectra are one type.● Only discrete λ are emitted.● Wavelengths of emitted light characteristic

of element.● Most interesting example was, and still is,

hydrogen.

Atomic Spectra

For hydrogen, it was discovered that the λ of the emitted light fit a simple equation:

22

21

111

nnRH

The Wave-Packet Model of Electromagnetic Radiation

● Hydrogen spectrum raised many difficult questions.

● Solution to the problem required thinking about light as particles.

● These particles are called photons.

The Wave-Packet Model of Electromagnetic Radiation

● The energy of a photon is related to the frequency of the corresponding light wave.

● E = hν● The constant of proportionality is called

Planck’s constant, h.h has units of energy × time.

The Wave-Packet Model of Electromagnetic Radiation

● Energy levels of an atom are quantized.● Emission occurs when an atom releases

energy in the form of light as a photon:

Atom(higher energy state) → Atom(lower energy state) + hν

The Bohr Model of the Atom

● Refinement of Rutherford Model.● Mass at center.● Electrons orbit the center like the moon

around the Earth.● The orbits are quantized.

The Bohr Model of the Atom

● Orbits● E is proportional to radius ● Quantized orbits● Absorption● Emission

The Bohr Model of the Atom

● Fantastic achievement.● Ushered in Quantum Mechanics.● Only worked for 1 e- systems.

The Energy States of the Hydrogen Atom

We already saw .111

22

21

nnRH

We now identify the two ni, n1 and n2, as labeling energy states of the hydrogen atom.

The Energy States of the Hydrogen Atom

Figure 3.5

Electromagnetic Radiation and Color

● Primary additive colorsRGB & color TVs

Figure 3.6

Electromagnetic Radiation and Color

● Primary subtractive colorsAbsorbed (subtracted) color determines the observed color in transmission

– A blue solution is absorbing yellow light (not blue light!).

Figure 3.6

The First Ionization Energy

● Ionization: removing an electron.● First Ionization: minimum amount of

energy needed to accomplish this.

The First Ionization Energy

Table 3.3

The First Ionization Energy

Figure 3.7

The Shell Model

● Data in Figure 3.7 support a shell model of electrons surrounding a nucleus.

● Core charge

The Shell Model and Periodic Table

● Connection between Figure 3.7 and Periodic table.

● Valence electrons● Core electrons

Photoelectron Spectroscopy (PES) and the Structure of Atoms

● PES uses light to ionize atoms, molecules or ions.

● Knowing the energy of the light, one can calculate the ionization energy

IE = hv – KE.

Photoelectron Spectroscopy and the Structure of Atoms

PES can remove an electron from any shell.Figure 3.14

Electron Configurations from PES

● PES reveals subshells.● n for shell.● s, p, d, f, … for subshell.● Superscripts for occupation number.

Electron Configurations from PES

Table 3.5

Allowed Combinations of Quantum Numbers

● Quantum numbersn

l

ml

Shells and Subshells of Orbitals

● The quantum numbers are used to label the shells and subshells.

● Three numbers are needed to specify an orbital.

● Each orbital can hold two electrons.

Orbitals and the Pauli Exclusion Principle

● There exists a fourth quantum number, ms.

● It can have one of two values: +½ or -½.● Each electron in an atom has a set of 4

numbers: (n, l, ml, ms).

● Pauli Exclusion Principle states that no two sets can be the same in a given atom.

Predicting Electron Configurations

Figure 3.23

Electron Configurations and the Periodic Table

● Relationship between electron configurations and periodic table.

Electron Configurations and Hund’s Rule

● No effect on electron configuration.● Important when drawing an orbital

diagram.Which of these is correct for carbon and why?

pss 221

pss 221

Electron Configurations and Hund’s Rule

● The left one adheres to Hund’s rule.

pss 221

pss 221

The Sizes of Atoms: Metallic Radii

Figure 3.25

The Sizes of Atoms: Covalent Radii

Figure 3.26

The Relative Sizes of Atoms and Their Ions

● Cations are smaller than the atoms from which they originate.

● Anions are larger than the atoms from which they originate.

Why?

Patterns in Ionic Radii

● Identical electron configurations come from isoelectronic species.

Second, Third, Fourth and Higher Ionization Energies

● Removal of an electron from a +1 cation is called the second ionization energy.

● Removal of an electron from a +2 cation is called the third ionization energy.

Average Valence Electron Energy

● AVEE● Weighted average.● Ionization energies used.

Average Valence Electron Energy

● Measure of attraction between electrons and nucleus.

● Measure of spacing of valence energy levels.

Average Valence Electron Energy

AVEE can be used to characterize a material as metal or nonmetal.

• Metal: low AVEE, small separation.

• Nonmetal: high AVEE, large separation.

AVEE and Metallicity

Figure 3.33