Chemistry 11 Resource: Chang’s Chemistry, Chapter 7.

THE ELECTRONIC STRUCTURE OF ATOMS

Chemistry 11

Resource: Chang’s Chemistry, Chapter 7

Objectives1. Explain how the lines in the emission spectrum

of hydrogen are related to the electron energy levels.

2. State the relative energies of s, p, d, and f orbitals in a single energy level.

3. State the maximum number of orbitals in a given energy level.

4. Draw the shape of s and p orbitals.5. Apply the Aufbau principle, Hund’s rule, and the

Pauli exclusion principle to write electronic configurations for atoms and ions up to Z = 20.

Activities Exercises from the text Quizzes 3-d models of atomic orbitals

Bohr’s model Ever since the 17th century, the

phenomenon of emission spectra has fascinated physicists.

The emission spectrum of a substance can be seen by energizing a sample of material.

Bohr’s model

Emission by a heated objectp 258p 267

Bohr’s model The emission spectra of gases are quite

different. Gases were found to emit light only at

certain wavelengths.

Bohr’s theory

Emission spectra of gasesp 268

Bohr’s theory What was the model of the atom before

Bohr? Could that model possibly explain the

emission spectra phenomenon?

Bohr’s theory Before Bohr, physicists knew that the

atom consisted of protons and electrons. They believed that the electrons moved

around the nucleus in circular orbits (Rutherford’s model).

Why was this model acceptable to scientists?

Bohr’s model In the early 20th century, Bohr added to

the contemporary model of the atom:

The single electron in the hydrogen atom can only be located in certain orbits.

Each orbit has a particular energy associated with it.

Bohr’s model

Bohr’s model of the atomp 269

Bohr’s model Only certain orbits are permitted. Each orbit has an associated energy value.

Therefore, the energy associated with e- motion is quantized, or fixed in value.

Bohr’s model Bohr attributed the emission spectrum of

hydrogen to the following process:The electron absorbs energy and jumps to a

higher orbit.When the electron returns to its ground

(normal) state, it emits energy through a photon (light particle).

Since only certain orbits (energy levels) are permitted, light at a certain wavelength is emitted.

Bohr’s model

Emission of light by a hydrogen atomp 269

Electron cloud model Bohr’s model could not account for the

emission spectra of atoms with more than one electron.

It became even more insufficient when physicists discovered that electrons are wavelike.

How can you pinpoint the location of an electron if it is a wave?

Electron cloud model Heisenberg’s uncertainty principle:

It is impossible to know [the momentum p and] the position of a particle with certainty.

How does this principle defy Bohr’s model of the atom?

Electron cloud model In the 1920s, Schrödinger applied this to the

model of the atom:

The exact location of an electron cannot be pinpointed.

Therefore, the representation of the electron was modified from lines to a cloud where an electron is more likely to be found.

Electron cloud model

The electron cloud modelp 278

Electron cloud model Schrödinger ushered in a new age of

physics called quantum mechanics. We now refer to the “location” of

electrons as atomic orbitals. Each atomic orbital has a certain

associated energy and a distribution of electron density.

Quantum numbers As a result of the discoveries in the 1920s,

electrons were assigned quantum numbers to describe their distribution or “location”.

Three quantum numbers are required to describe the distribution of electrons.the principal quantum number nthe angular momentum quantum number lthe magnetic quantum number ml

Quantum numbers The principal quantum number n is

designated an integer value greater than 0, i.e. 1, 2, 3, 4, …

It relates to the average distance of the e- from the nucleus.

The larger n is, the farther away it is from the nucleus.

If n is larger, is the orbital bigger or smaller?

Quantum numbers The angular momentum quantum number

l tells us the “shape” of the orbital. l is related to n

The values of l can vary from 0 to (n -1).

If n = 1, what are the possible values of l? What if n = 3?

Quantum numbers The value of l is generally designated by

the letters s, p, d, … as follows:

l 0 1 2 3 4 5Name of orbital s p d f g h

Quantum numbers

If an e- has a principal quantum number of 1 (n = 1), how many orbitals are possible?

Quantum numbersSince n = 1, the only possible value of l is 0.

remember:l varies from 0 to n – 1

since n – 1 = 0,0 is the only possible l value

therefore:there is only 1 orbital when n = 1.

This is called the 1s orbital.

l 0 1 2 3 4 5

Name of orbital s p d f g h

Quantum numbers

If an e- has a principal quantum number of 2, how many orbitals are possible?

Quantum numbers

If n = 2, l can be 0 and 1

therefore:TWO orbitals are possible.

These orbitals are called 2s and 2p.

l 0 1 2 3 4 5

Name of orbital s p d f g h

Quantum numbers A group of orbitals that have the same

value for n (e.g. 2s and 2p) are frequently called a shell.

Quantum numbers

The magnetic quantum number ml describes the orbital’s orientation in space.

The value of ml depends on l and varies as follows:

-l, (-l +1), … 0, … (l - 1), l

Quantum numbers

If n = 2 and l = 1, how many orbitals are possible?

Quantum numbersThree orbitals in that subshell are possible:

since l = 1,ml = -1, 0, 1

Therefore:3 orbitals are possible.

These orbitals are called 2px, 2py, and 2pz.

This will all make a little more sense later on

Quantum numbers

Relation between quantum numbers and atomic orbitals

n l mlNumber of

orbitalsAtomic orbital designations

1 0 0 1 1s

2

3

Quantum numbers

Relation between quantum numbers and atomic orbitals

n l mlNumber of

orbitalsAtomic orbital designations

1 0 0 1 1s

201

0-1, 0, -1

13

2s2px, 2py, 2pz

3012

0-1, 0, -1

-2, -1, 0, 1, 2

135

3s3px, 3py,, 3pz

3dxy, 3dyz, 3dxz, 3dx2-

y2, 3dz2

Quantum numbers

A fourth quantum number ms is used to denote the spin of the electron.

Electrons are known to spin two ways: up or down.

This electron spin quantum number will be discussed later on.

Atomic orbitals Both Bohr and Schrödinger made

significant contributions to our understanding of the atom.

We will use their ideas to get a better picture of atomic structure.

Atomic orbitals In principle, an electron can be found

anywhere in the atom.

In a typical hydrogen atom, where would the single electron most likely be?

Atomic orbitals Common sense dictates that the single

electron will probably be close to the nucleus.

Thus we can represent the 1s orbital by drawing a boundary that encloses about 90% of the total electron density:

p 282

Atomic orbitals Recall that each value of n has an s orbital

(1s, 2s, 3s, …)

The shape of the s orbitalp 282

How does the value of n affect the shape/size of the orbital?

Atomic orbitals

At what value for n do we see s orbitals?

Atomic orbitals

There is an s orbital at every value of n.Think of it as the “basic” orbital.

Atomic orbitals

If n = 1, does a p orbital (l = 1) exist?

Atomic orbitals

No.p orbitals exist when n = 2 or higher:

when n = 1, l = 0;therefore only the 1s is possible.

p orbitals are associated with l = 1.

Atomic orbitals p orbitals appear when n is 2 or higher:

since n = 2,l = 0, 1

These correspond to the 2s and the 2p orbitals.

Atomic orbitals Furthermore, when n = 2 and l = 1 (2p)

ml = -1, 0, 1

Therefore there are THREE possible 2p orbitals.

Atomic orbitals

Shape of the 2p orbitalsp 283

Atomic orbitals

The shape of the 3d orbitalsp 283

Atomic orbitals Remember that each orbital has a shape

(cloud) and a certain energy associated with it.

Which orbital has the lowest energy associated with it?

Atomic orbitals

The 1s orbital has the lowest energy.

Atomic orbitals

Orbital energy levelsp 285

Atomic orbitals

The order in which atomic subshells are filled

p 285

Electron configuration An electron can be identified by its four

quantum numbers. You may think of the quantum numbers as

the “address” of the e- because they describe its location.

Electron configuration

Summary of quantum numbers

Quantum number

Information about e-Part of the “address”

n distance from nucleus province / state

l shape city

ml orientation in space street

ms spin number

Electron configuration

What are the four quantum numbers of hydrogen’s single electron?

(n, l, ml, ms)

Electron configuration

(n, l, ml, ms)

(1, 0, 0, +½)or

(1, 0, 0, -½)

Electron configuration

Write the four quantum numbers of an electron in the 3p orbital.

Electron configuration Homework

p 299: 7.55-7.61, odd; 7.64 Review for a quiz next class

Electron configuration The electron configuration of an atom is

how the electrons are distributed among the various atomic orbitals.

This is the electron configuration of hydrogen which has 1 e-.

1s1Denotes the principal quantum number n

Denotes the angular momentum quantum number l

Denotes the number of electrons in the orbital or subshell

Electron configuration Electron configuration can also be

represented by an orbital diagram that shows the spin of the electron:

1s1

Pauli exclusion principle No two electrons in an atom can have the

same four quantum numbers. If they are in the same orbital (i.e. same

values for n, l, and ml) then they must have different values for ms.

1s21s2 1s2

Hund’s rule The most stable arrangement of electrons

in subshells is the one with the greatest number of parallel spins (same spins).

Hund’s rule Carbon (Z = 6) is 1s22s22p2

Which configuration satisfies Hund’s rule?

1s2 2p22s2

1s2 2p22s2

1s2 2p22s2

Aufbau principle “Aufbau” is the German word for “building

up” Just as protons are added one-by-one to

build up the elements, so are electrons into the atomic orbitals.

This introduces a different way of showing of electronic configuration.

Aufbau principle The configuration shows the noble gas (in

brackets) that most nearly precedes the element being considered.

So instead of:Na 1s22s22p63s1

You may represent Na as:Na [Ne]3s1

Very convenient.

The transition metals The electronic configurations of elements

from Z = 1 to Z = 20 are relatively straightforward.

The electronic configurations of the transition metals have “strange” electronic configurations that do not necessarily follow convention (p292)

Why do you think this is?

Electronic configuration Homework

pp 299 – 302○ # 7.71, 72, 79. 85. 87. 124

Quiz next class Bring old newspapers

Atomic orbital models You will divide yourselves into THREE

groups of fairly equal numbers. Each group has a different assignment. This will count as a project. It is due on 27 January 2009.

Atomic orbital models Group 1

Build models for the 1s, 2s, and 3s orbitals.Create 3 posters:○ Orbital energy levels○ Order in which atomic subshells are filled○ Pauli exclusion principle and Hund’s rule

Group 2Build models for the 2p and 3p orbitals.

Group 3Build models for the 3d orbitals.

Atomic orbital models Color codes:

Blue: n = 1Yellow: n = 2Red: n = 3

Atomic orbital models

Grading the orbital models (for each model)

Correct shape 5 pts

Correct size(relative to similar orbitals with different values for n)

5

Correct color 2

Stability 3

Total 15

Atomic orbital models

Grading the posters (for each poster)

Accurate information 5 pts

Pertinent details present 5

Organization of information 3

Creativity / aesthetic value 2

Total 15