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Modern Atomic Theory and

Electron Configurations

Lectures 8 and 9

Chem 101

Types of Electromagnetic Radiation

Electromagnetic radiation is given off by atoms when they have been

excited by any form of energy, as shown in flame tests.

Electromagnetic radiation

(a beam of light) can be pictured in two ways:

as a wave and as a stream of individual protons.

The wavelength of a

wave is the distance between peaks.

A photon of red light

(relatively long wavelength) carries less energy than

does a photon of blue light (relatively short wavelength).

Properties of Electromagnetic Waves

• Velocity = c = speed of light – 2.997925 x 108 m/s ( use 3.00 x 108 m/s ) – All types of light energy travel at the same speed.

• Amplitude = A = measure of the intensity of the wave, i.e.“brightness”

• Wavelength = = distance between two consecutive peaks or troughs in

a wave – Generally measured in nanometers (1 nm = 10-9 m)

• Frequency = = the number of waves that pass a point in space in one

second – Generally measured in Hertz (Hz), – 1 Hz = 1 wave/sec = 1 sec-1

• c =

• Energy = h ν Energy is equal to Planck’s constant times frequency – Planck’s constant, h, is 6.63 x 10-34 Joule seconds

When salts containing Li+, Cu2+, and Na+

dissolved in methyl alcohol are set on fire, brilliant colors result:

Li+, red; Cu2+, green; and Na+, yellow.

Hmco Photo Files

Emission of Energy by

Atoms/Atomic Spectra

• Atoms that have gained extra energy release that

energy in the form of light.

Atomic Spectra

• Line spectrum: very specific wavelengths

of light that atoms give off or gain

• Each element has its own line spectrum,

which can be used to identify that element.

When an excited H atom returns to a lower energy

level, it emits a photon that contains the energy released by the atom.

Each photon emitted by an excited hydrogen atom corresponds

to a particular energy change in the hydrogen

The colors and wavelengths

(in nanometers) of the photons in the visible

region that are emitted by excited hydrogen atoms.

Atomic Spectra

• The atom is quantized, i.e. only certain energies

are allowed.

The Bohr model of the hydrogen atom represented the electron

as restricted to certain circular orbits around the nucleus.

Energy of electron is

related to the distance of

electron from the nucleus

Bohr’s Model

• Energy of the atom is quantized – Atom can only have certain specific energy states called

quantum levels or energy levels.

– When atom gains energy, electron “moves” to a higher

quantum level

– When atom loses energy, electron “moves” to a lower

energy level

– Lines in spectrum correspond to the

difference in energy between levels

(a) The hydrogen 1s orbital. (b) The size of the orbital is defined by a sphere that

contains 90% of the total electron probability.

Bohr’s Model

• Ground state: minimum energy of an atom – Therefore electrons do

not crash into the nucleus

• The ground state of hydrogen corresponds to having its one electron in the n=1 level

• Excited states: energy levels higher than the ground state

Orbitals and Energy Levels

• Valence shell: the highest-energy occupied ground state orbit

• Regions in space of high probability for finding the electron. These are called orbitals.

• Each principal energy level contains one or more sublevels. Sublevels are made up of orbitals.

• Each type of sublevel has a different shape each and energy.

• Each sublevel contains one or more orbitals.

The 1s orbital.

A diagram of principal energy levels 1 and

2 showing the shapes of orbitals that compose the sublevels.

The shapes

and labels of the five 3d orbitals.

Pauli Exclusion Principle

• No orbital may have more than 2 electrons.

• Electrons in the same orbital must have opposite

spins.

• s sublevel holds 2 electrons (1 orbital)

• p sublevel holds 6 electrons (3 orbitals)

• d sublevel holds 10 electrons (5 orbitals)

• f sublevel holds 14 electrons (7 orbitals)

• For a multiple-electron atom, build-up the energy

levels, filling each orbital in succession by energy

• Degenerate orbitals: orbitals with the same

energy

– e.g. Each p sublevel has 3 degenerate p orbitals

Orbitals, Sublevels & Electrons

Electron Configurations

• For a set of degenerate orbitals, fill each orbital half-

way first before pairing

• Electron configurations show how many electrons

are in each sublevel of an atom – describes where

electrons are.

- 1s22s1 is the electron configuration for a ground

state Li

- 1s22s22p3 is for nitrogen

Electron Configurations

• Valence shell: highest energy level

– Electrons in the valence shell are called valence electrons.

– Core electrons: electrons not in the valence shell

– Often use symbol of previous noble gas in brackets to represent core electrons, giving

[He]2s22p3 for nitrogen or [Ne]3s2 for magnesium

Electron Configuration

and the Periodic Table

• Elements in the same column on the

periodic table have:

– Similar chemical and physical properties

– Similar valence shell electron configurations

• same numbers of valence electrons

• same orbital types

• different energy levels

The electron configurations in the sublevel last occupied for

the first eighteen elements.

s1

s2

d1 d2 d

3 d4 d5 d6 d7 d8 d9 d10

p1 p2 p3 p4 p5 s2

p6

f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14

1

2

3

4

5

6

7

The orbitals being filled for elements in various parts

of the periodic table.

A box diagram showing the

order in which orbitals fill to produce the atoms

in the periodic table. Each box can hold two electrons.

The periodic table with atomic symbols, atomic numbers, and

partial electron configurations.

Metallic Character:

– Form cations

– Lose electrons in reactions – oxidized

– Oxidation is Loss of electrons - OIL

– The easier it is for an element to lose

electrons, the more metallic character

is has.

The classification of

elements as metals, nonmetals, and metalloids.

Metallic Character

• Reactivity of metals increases to the left on the

period and down in the column

– Follows ease of losing an electron

• Reactivity of nonmetals (excluding the noble

gases) increases to the right on the period and up

in the column

Trend in Ionization Energy

• Minimum energy needed to remove a valence electron from an atom

– Gas state

• The lower the ionization energy, the easier it is to remove the electron.

– Metals have low ionization energies

• Ionization energy decreases down the group.

– Valence electron farther from nucleus

• Ionization energy increases across the period.

– Left to right

The Group 1 elements: the farther down a group and element

appears, the more likely it is to lose an electron.

The Group 2 elements: the farther down a group and element

appears, the more likely it is to lose an electron.

Ionization energies tend to decrease in

going from the top to the bottom of a group.

Ionization energies tend to increase from left to right

across a given period on the periodic table.

Relative atomic sizes for selected atoms. Note that atomic size increases down a

group and decreases across a period.

Electronegativity

40

• Measure of the ability of an atom to attract

shared electrons

• Larger electronegativity means atom

attracts shared electrons more strongly