Download Modern Atomic Theory

Chem 101
Modern Atomic Theory
and
Electron Configurations
Lectures 8 and 9
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)
Orbitals, Sublevels & Electrons
• 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
Electron Configurations
• For a set of degenerate orbitals, fill each orbital halfway 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
1
2
3
4
5
6
7
p 1 p 2 p 3 p 4 p 5 s2
p6
s2
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
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
• Measure of the ability of an atom to attract
shared electrons
• Larger electronegativity means atom
attracts shared electrons more strongly
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