1 Atomic Structure Electron Configurations & Periodicity.

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Atomic StructureAtomic StructureElectron Configurations & Electron Configurations & PeriodicityPeriodicity

Introduction Atomic structure explains chemical properties

and patterns of chemical reactivity. Chemical reactions involve electrons. Knowing

where the electrons are, how many, and what their energy levels helps explain many chemical phenomena.

Spectroscopy is used to explore atomic structure. Because of this, we start with a discussion of the nature of light.

The Nature of LightParticles or Waves??

Light must be made of particles because it… travels in a vacuum reflects off of objects exerts force (on the

tails of comets) Light must consist of

waves because it… reflects like waves refracts and diffracts exhibits interference

By the end of the 19th century,scientists had concluded that light is composed of WAVES!

Electromagnetic Radiation Visible light is just one form of electromagnetic

radiation Light propagates in space as a wavewave In vacuum, speed of light is constant and given

the symbol “c”, c = 3.00 x 108 m/s Light waves have amplitude, frequency, and

wavelength. = wavelength - distance between consecutive

crests, = frequency, # of crests that pass a given point in one second (SI Unit is s-1 or Hz)

c = c = and = c/

James Maxwell proposed in the mid-1800’s that light is composed of perpendicular electric and magnetic waves

are distances – S.I. unit is the meter

Which of the above waves has the higher frequency?

The 7 Regions of the E.M. Spectrum

Sample Calculation – wavelength/frequency conversion

Calculate the frequency of visible light having a wavelength of 485 nm?

Remember to use S.I. units in your calculations!

= c / = (3.00 x 108 m/s) ÷ (4.85 x 10-7 m) = 6.19 x 1014 s-1

What colour of visible light is this?

Sample Calculation – wavelength & frequency conversion

CO2 absorbs light with a wavelength of 0.018 mm. Which frequency is this?

Remember that 1 Hertz (Hz) 1 s-1

What is the wavelength of WFAE at 90.7 MHz?

8 113 1

3

c 3.00 10 ms1.7 10 s

0.018 10 m

8 1

6 1

c 3.00 10 ms3.31m

90.7 10 s

Particle Theory of Light In 1900, Max Planck turned the world of physics on

its head by resurrecting the particle theory of light. Planck was trying to explain the “ultraviolet

catastrophe” while studying blackbody radiation. He proposed that light is composed of particles

(quanta) each carrying a fixed amount of energy. The amount of energy per quantum is directly

proportional to the frequency of the light:

34hcE h or E , h 6.626 10 J s

34hc

E h or E , h 6.626 10 J s

Planck’s Equation – sample problem The wavelength of maximum visual acuity in

humans is 550 nm. (green light) What is the energy of a single photon having this

wavelength?

Ans. 3.61 x 10-19 J per photon

What is the energy of a mole of photons at 550 nm?

Ans. 218 kJ/mol

Energy, wavelength, frequency

a) X-Rays

b) Red light

c) Green light

d) Radio waves

Rank the above in order of … increasing wavelength decreasing energy increasing frequency

Atomic Emission Spectra“Line Spectra” vs “Continuous Spectra”

Historic work (1800’s) involving light emitted by pure elements in gas phase, subjected to very high voltages.

The light was passed through a prism or diffraction grating to produce a spectrum.

Each element emitted only certain wavelengths of light – a LINE SPECTRUM instead of the more familiar “continuous” spectrum seen in the rainbow.

Balmer and Rydberg described the lines mathematically, for hydrogen.

Hydrogen’s emission spectrum has several lines in the visible region – called the “Balmer series”

A Continuous Spectrumcontains all wavelengths of light – a rainbow

Bohr’s Explanation of Line Spectra

Neils Bohr developed a mathematical model that could explain the observation of atomic line spectra.

He proposed that electrons orbit the nucleus in certain “allowed” orbits - or energy levels. That is, the electron’s energy is QUANTIZED (not continuous).

Working on a model of single-electron atoms, Bohr derived an equation to calculate the energy of an electron in the nth orbit of such an atom:

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22.18 10n

ZE

n

where Z = atomic number

and n = electron energy level

Bohr’s Quantum Model for Hydrogen

The electron in hydrogen occupies discrete energy levels.

The atom does not radiate energy when the electron is in an energy level.

When an electron falls to a lower energy level, a quantum of radiation is released with energy equal to the difference between energy levels.

An electron can jump to higher energy levels if the atom absorbs a quantum of radiation with sufficient energy.

The Hydrogen Line Spectrum

Bohr’s Model: Calculations

Show that an electron transition from n = 4 to n = 2 results in the emission of visible light. Calculate the wavelength of the light emitted.

Step 1: Calculate E4 & E2

E4 = -1.36 x 10-19 J

E2 = - 5.45 x 10-19 J

Step 2: Find E42

E = E2 – E4 = 4.09 x 10-19 J

Step 3: Calculate = 4.86 x 10-7 m = 486 nm

What’s Wrong with Bohr’s Model??

Although it works great for single-electron atoms, Bohr’s model fails for atoms with 2 or more electrons!

It was a huge leap forward, but was fundamentally flawed.

Ultimately, the failure of Bohr’s model lay in the fact that he treated the electron as a charged particle orbiting the nucleus like a planet around the sun.

Electrons are more complicated …

Wave-Particle Dual Nature of Electrons

Einstein’s famous equation, E = mc2, suggests that energy and mass are related – one can convert matter directly into energy.

Louis de Broglie made the leap that if light can behave as wave/particles, then so can matter!

He showed that the wavelength of a baseball is negligible (as expected), but that the wavelength of an electron was on the order of magnitude equal to that of electromagnetic radiation!

This is what Bohr had missed – an atomic model must make use of the wave-nature of electrons to be complete!

Visualizing the Wave-Nature of Electrons

Heisenberg’s “Uncertainty Principle” The more precisely the position is

determined, the less precisely the momentum is known in this

instant, and vice versa.

It is impossible to determine BOTH the position and momentum (velocity) of an electron - the act of observing the electron changes it.

The electron exists in nature as a wave-particle (whatever that is) until an experimenter observes it and imposes one of the two natures upon it!

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"When it comes to atoms, language can only be used as in poetry. The poet too, is not nearly so concerned with describing facts as with creating images."

- Niels Bohr to Werner Heisenberg

The Heisenburg uncertainty principle is the catch 22 of physics. Most modern physics and chemistry is largely based on the study of electrons. In order to view an electron we must use light to observe it, but in the process of observing we change the electron’s momentary existence because the light necessary to view it takes the form of a photon particle. The photon particle travels at approximately the same velocity as the electron particle, so that when it impacts the electron it alters its position as when two asteroids collide into each other at similar velocities but different angles.

Therefore, according to Heisenburg, it is impossible to know decisively an electron’s position and momentum.

Heisenburg proposed that this inability to perceive reality on an atomic level is not the result of technological ignorance, but rather a fact of the laws of nature. In other words, Heisenburg believed that no matter how technologically advanced we become we would never be able to fully perceive and record atomic reality. This means that an electron’s orbital path around the nucleus of an atom can only be statistically approximated.

This is why the quantum atomic model has a cloud of electrons, not a few electrons following exact elliptical orbits as in Bohr's atomic model. Each electron in the quantum atomic model is a calculated probability, so the cloud describes not the exact orbit of millions of electrons, but rather the probability of the position of only a few electrons at any given time.

The Schrödinger Wave Equation Schrödinger applied de Broglie’s concept of

matter waves to the electron He explained the quantum energies of the

electron orbits in the Bohr model of the atom as vibration frequencies of electron "matter waves" around the atom's nucleus.

He provided a mathematical equation (wave function) to describe electron waves.

Physicists had difficulty explaining the MEANING of the wave function itself, but it turns out that its square gives the probabilityprobability of finding the electron in a particular region of space in the atom.

Schrodinger Equation, cont. These regions of probability are called orbitals, orbitals, and this

concept replaces that of the electron “orbit”. Instead of electrons orbiting like planets around the sun,

Schrodinger’s picture shows an “electron cloud” surrounding the nucleus and doesn’t state anything about the path or position of electrons within that cloud.

Orbital “pictures” shown in texts represent regions of space where the probability of finding an electron is 90% or greater.

Quantum Numbers Schrodinger's wave equation used three constants

(“quantum numbers”) that were used to describe orbitals – regions of space where an electron has a probability of being found.

His equation used 3 quantum numbersquantum numbers to describe the size & energy, shape, and orientations of the orbitals in atoms.

The three quantum numbers: n, l and ml . This set of three numbers provides us with an “orbital

address” for an electron within an atom – we can describe the orbital in which an electron is found if we know these three quantum numbers.

Principal Quantum Number, n

n can have values > 1, integers only. For element atoms on the periodic table in their

ground statesground states, 1 n 6 As n increases, energy of electron increases and

size of orbital increases Larger n means greater probability of greater

distance from the nucleus. Electrons with same value of n are said to be in

the same “electron shell”

Shape Quantum Number, l The values for l (lower case letter L, script) are limited by

the principal energy level – they range from 0 to “n–1” Example: If n = 3, l can be 0, 1 or 2

l gives information about shape of orbital

l = 0 s orbital (found on all principle energy levels)

l = 1 p orbitals (found on level 2 and higher)

l = 2 d orbitals (found on level 3 and higher)

l = 3 f orbitals (found on level 4 and higher)

s Orbitals on three energy levels

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Graphs show Probability vs Distance from nucleus

Graphs show Probability vs Distance from nucleus

p and d Orbital shapes

Orientation Quantum Number, ml The values of ml are limited by the value of l for

a given orbital. They include positive and negative integers from -l through +l , inclusive

For a p orbital, l = 1 , so ml can be: 1, 0, 1

ml gives information about orientationorientation of orbitals in space

The # of allowed values gives the number of orbitals of this type on a given energy level. (3, in the case of p orbitals)

Electron Spin Quantum Number, ms A fourth quantum number, ms

describes electron spinelectron spin (either +½ or –½)

Each electron in atom has a unique set of these four quantum numbers.

Electrons in orbitals with same n and l values are said to be in the same subshell.subshell.

Electrons with all three numbers the same, n, l , and ml , are in

the same orbital.

A spinning negative charge creates a magnetic field. The direction of spin d determines the direction of the field.

Pauli Exclusion Principle Electrons have negative charge and repel each

other. How are the electrons in an atom distributed?

Wolfgang Pauli proposed that no two electrons no two electrons in a given atom can be described by the in a given atom can be described by the same four quantum numbers!same four quantum numbers!

The first three quantum numbers determine an orbital – therefore spins must be opposite!

Practical resultPractical result is that each orbital can hold a maximum of two electrons, with opposite spins.

Concept Check Which of the following sets of

quantum numbers are NOT allowed in the hydrogen atoms? For those that are incorrect, state what is wrong.

n l mlms

2 1 -1 + ½

1 1 0 - ½

8 7 -6 + ½

1 0 2 - ½

What is the maximum number of electrons in an atom that can have these quantum numbers?

n = 4 n = 5, ml = +1

n = 5, ms = + ½

n = 2, l = 1 n = 3, l = 2 n = 3, ml = 4

Electron Configurations

Electrons orbitals are defined by their quantum numbers, n, l and ml .

Each electron in an atom has a unique set of 4 quantum numbers.

No two electrons can have the same “address”, i.e., the same 4 quantum #’s

Rules define how multiple electrons will be distributed among the possible energy levels

Aufbau Principle In the ground state, the electrons occupy the lowest

available energy levels. An atom is in an excited state if one or more electrons are in higher energy orbitals.

In atoms with more than one electron the lower energy orbitals get filled by electrons first!

This is the Aufbau principle, which is named after the German word which means “to build up”.

When one describes the locations of the electrons in an atom, start with the lowest energy electron and work up to the highest energy electron.

Hund’s Rule A set of orbitals is said to be “degenerate” if the orbitals

possess the same energies. For example, all three “2p” orbitals on energy level 2 are degenerate. All five “3d” orbitals on energy level 3 are also degenerate.

When filling a set of degenerate orbitals, Hund said that electrons should be left unpaired as long as possible so as to minimize electron-electron repulsions within the orbitals.

When each degenerate orbital has one electron, electrons will then pair, spins opposed, until that sub-shell is filled.

Summary of Distribution Rules Electrons distribute to lower energy levels

until they are filled, before occupying higher levels. (Aufbau Principle)

Electrons will spread out as much as possible within a sub-shell. (Hund’s Rule)

Electrons will pair up, two to an orbital, spins opposed, until that sub-shell is filled. (Pauli Exclusion)

Every electron will have unique set of 4 quantum numbers.

Electronic Configurations An electron configuration is a way of describing the

locations of electrons within an atom. Usually written for the ground state of an atom: when its

electrons occupy orbitals giving the lowest possible overall energy for the atom

Each subshell is designated by “n” and the type of orbital. Such as: 1s 3p 4d

The number of electrons in an orbital is shown as a superscript: 1s2 3p3 4d9

How does one determine the order of filling of the orbitals in an atom?? Which orbitals get filled first??

Orbital Energies in Multielectron Atoms

The Order of Filling of OrbitalsSimply follow the order of elements in the periodic table!

Another useful device…

Start at the top and follow the arrows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s…

Examples of Electron Configurations Helium has 2 electrons in the 1s orbital, He: 1s2

Carbon has 6 electrons, C: 1s2 2s2 2p2

Calcium has 20 electrons, Ca: 1s2 2s2 2p6 3s2 3p6 4s2

Calcium cation, Ca2+: 1s2 2s2 2p6 3s2 3p6 Sulfide anion, S2–: 1s2 2s2 2p6 3s2 3p6

Notice that the superscripts add up to the total number of electrons on the atom or ion! Each of these examples represents the “ground state” of the atom/ion because the electrons are in the lowest possible energy levels.

Orbital Diagrams

Noble Gas “shorthand” Notation The core electrons can be indicated by a short

form: use the last noble gas to indicate filled shells.

Find the last noble gas before the element under consideration. Start the electron configuration with the symbol for this noble gas element and just tack on the extra electrons:

Ca: 1s2 2s2 2p6 3s2 3p6 4s2

Ca: [Ar] 4s2

Ar: 1s2 2s2 2p6 3s2 3p6

There Have to be EXCEPTIONS As atoms get larger, their

outermost electrons become closer in energy

Large atoms have “exceptions” to the Aufbau Principle in the locations of their valence electrons. We will memorize two: Cr and Cu

Transition metal cations are also strange in their electron configurations

Cr: [Ar] 4s1 3d5

Cu: [Ar] 4s1 3d10

In both cases, one of the 4s electrons is promoted to the “3d” subshell.

Fe2+: [Ar] 3d6

Cu2+: [Ar] 3d9

Notice that the 4s electrons are lost before the 3d electrons!

Paramagnetism & Diamagnetism An atom is said to be

paramagnetic if it possesses unpaired electrons.

These electrons will result in the atom possessing a net magnetic field

It will then be attracted into an external magnetic field

An atom is diamagnetic if all its electrons are paired – it will NOT possess a net magnetic field because each electron’s field will be cancelled

These atoms are NOT attracted into an external magnetic field

What is “ferromagnetism”?

Testing for ParamagnetismSee Kotz & Treichel page 290

Electron ConfigurationCauses Periodic Trends “Periodicity” refers to similarities in behavior and

reactivity due to similar outer shell electron configurations.

All the Alkali Metals have one unpaired valence electron; all the Noble Gases have completely filled subshells

We will examine periodic trends in atomic radius, ionization energy, electronegativity, and electron affinity

“Valence Electrons”

Valence electrons are those electrons in the highest principal energy level within an atom - the s and p electrons in the outermost shell of an atom in its ground state.

These are the electrons involved in forming bonds with other atoms during chemical reactions.

Most common oxidation states (ion charge) for the element can be derived from valence electrons

Completely filled, half filled and empty sub-shells have special stability (not sure why!?)

Valence Configurations

The “valence configuration” is that part of an electron configuration that describes the valence electrons.

For example, sodium’s valence configuration is just “3s1”.

The valence configuration of bromine is “4s24d5”. We leave out the “3d10” electrons because they are not in the outermost principal energy level.

Periodic Trends: Atomic Size Properties of elements and their ions are

determined by electron configurations Individual atoms cannot be accurately

measured. Atomic sizes are determined from bond

lengths of compounds with various atoms

Br-Br2.28 Ar = 1.14 A

Br-C1.91 Arcarbon = 1.91-1.14 = 0.77 A

Atomic Size, cont. Atomic radii increase within a group (column) as

the principal quantum number of outermost shell increases

Electrons of outermost shells have greater probability of being farther from the nucleus

Atomic size decreases across a row (period) from Left to Right, because the effective nuclear charge increases. There are more and more protons in the nuclei

attracting more and more electrons!

Atomic Radii:Atomic Radii:

Atoms get larger as you move down a group

Atoms get smaller as you go across a period

Ionization Energies First Ionization energy: energy required to

remove an electron from the ground state atom in the gaseous state to make a cation

1st ionization of Mg:

Mg Mg+ I.E. = 738 J/mol

[Ne]3s2 [Ne]3s1

2nd ionization of Mg

Mg + Mg2+ 1450 J/mol

[Ne]3s1 [Ne]

Trend in First Ionization Energies First ionization energies decrease down a group

as radii increase: Na > K > Cs First ionization energies increase across a row

as radii decrease: Li < F < Ne For similar reasons that explained trends in

atomic radii. For a given element, IE2 always > IE1 because of

the same # protons in nucleus holding onto fewer electrons.

Trends in First Ionization Energies

First Ionization Energy (cont’d) Lowest IE1 are in the

lower left corner of periodic table (Fr and Cs)

Irregularities in IE1 trends arise from special electron configurations.

Half filled sub-shells are especially stable: every orbital has only one electron.

Oxygen has slightly lower IE1 than nitrogen because it takes more energy to disrupt N’s half-filled orbital configuration

Boron has slightly lower IE1 than Beryllium because its first electron is removed from the p-subshell further from the nucleus than the first electron removed in Be

Electron Affinity Energy change when a gas phase atom gains an

electron to become an anion.

Cl (g) + e- Cl– (g) EA = – 349 kJ/mol

The more negative an EA, the greater the affinity for electrons!

Electron affinities “tend” to become more negative across a period and less negative moving down a group

EA least negative for alkaline earth and noble gas elements: they have full or half full s- or p-orbitals and it would take a lot of energy to add another electron

Ion Sizes Cations: radius is always smaller than the

neutral atom. Loss of electrons (to make cation) leaves fewer

electrons attracted to same # of protons OR….. loss of outermost shell altogether as in

group 1A [Ne]3s1 [Ne] for sodium Anions: radius is always larger than neutral

atom. Same # protons pulling on more electrons;

also electrons repel each other.

Isoelectronic Atoms and Ions Isoelectronic atoms or ions have same electron configurations – the same number of electrons in same orbitals

Na+ , Ne and F- are isoelectronic: 1s2 2s2 2p6

For isoelectronic atoms and ions, the species with the most protons will be smallest

It has greater nuclear attraction for the same number of electrons, so the electron shells will be pulled closer to nucleus.

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Bonding in Molecules

Ionic Bonds

Polar Covalent Bonds

Non-Polar Covalent Bonds

Definitions

ELECTRONEGATIVITY is the attraction an atom has for electrons in a chemical bond.

A POLAR BOND is one between atoms with different electronegativities - one atom attracts electrons more strongly than the other. As a result, one atom acquires a partial positive charge and the other a partial negative charge.

More Definitions

A DIPOLE is a separation of charge. A polar bond is an example of a dipole.

A DIPOLE MOMENT is the magnitude of a dipole - it depends on the size of the charges involved.

A POLAR MOLECULE is one that possesses an overall dipole moment.

Classifying Chemical Bonds

IONIC bonds are characterized as… bonds between metals and non-metals bonds between cations and anions bonds involving a transfer of electrons (from a

metal to a non-metal)

The best classification of an ionic bond uses the concept of electronegativity.


An ionic bond forms between two atoms with a large difference in electronegativity - usually stated as being greater than about 1.7 Paulings.

e.g. NaCl: difference is 2.0 Paulings e.g. CuO: difference is 1.7 Paulings


COVALENT bonds are usually described as… bonds between two non-metal atoms bonds involving a sharing of electrons

Using the concept of electronegativity, we can see that there are two types of covalent bonds.


POLAR COVALENT BONDS result from unequal sharing of electrons by atoms that creates a small dipole.

This type of bond occurs when there is a small but significant difference in electronegativity - between 0.4 and 1.7 Paulings. e.g. N-H: diff is 0.9 Paulings e.g. Cu-S: diff is 0.7 Paulings


NON-POLAR COVALENT BONDS result from an equal sharing of electrons.

There is no dipole created since the electrons are shared equally.

This bond forms between atoms with similar (or the same) electronegativities - differences between 0 and 0.4 Paulings.

Exercise

Classify the following chemical bonds. Show your work. Li-F Br-Br Fe-O Al-Cl C-S C-H

Polarity of Molecules

A diatomic molecule (e.g. HF) will be polar if its bond is polar.

Larger molecules are polar only when the following criteria are met: they possess at least one polar bond their geometries do not result in the dipoles

cancelling each other out! (e.g. CO2)

1 Atomic Structure Electron Configurations & Periodicity.

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