Electron ConfigurationsChapter 4

Chemistry-CP

Radiant Energy

• The understanding of how electrons behave comes from studies of how light interacts with matter.

• Light carries energy through space in the form of waves and also in the form of extremely tiny, fast moving particles.– Light has the properties of waves & particles.

Light as Waves

• Light waves are a form of electromagnetic radiation

Waves• Light travels at the speed of light.

– The speed of light is constant, which means it is always the same value: 3.00 × 108 m/s

• Because light moves at a constant speed, wavelength & frequency are inversely proportional as per the following equation.

c = ×

c = speed of light = wavelength (lambda) = frequency (measured in 1/s or s-1 or

Hertz (Hz))

c = × • What is the frequency of a wave having a wavelength of

8.12 x 102 m?

• A helium neon laser produces red light whose wavelength is 633 nm. What is the frequency of this radiation?

• Calculate the wavelength of a radio wave with a frequency of 9.31 × 106 s-1.

Gamma Rays

• Generated by radioactive atoms, nuclear explosions and supernova explosions

• Can kill living cells—used for cancer treatment

• Used to sterilize medical equipment•

http://www.youtube.com/watch?v=NZF3_e6_xj4

X-Rays• Discovered by accident in 1895, when W.C. Roentgen

shielded a cathode ray tube with black paper and found that a fluorescent light could be seen on a screen a few feet from the tube (first bone x-ray was of his wife’s hand!)

• Electrons shot at an element (such as tungsten or molybdenum) with high energy can knock an electron out of that atom, producing x-rays

• Used for radiography, crystallography, astronomy, airport security

Ultraviolet Radiation

• Gets its name from the fact that it consists of waves with frequencies higher than what humans associate with violet light

• Emitted from the sun, from black lights• UV-B produces Vitamin D, too much = DNA

damage & collagen fibers, can cause sunburn, may lead to cataracts

• Some animals, insects, birds and reptiles can see the near ultraviolet making certain flowers, etc. brigher to them.

• Portion of the electromagnetic spectrum that is visible to the human eye

• ROYGBIV—Violet has the highest frequency

Infrared • “Below red”• “Heat radiation”• Emitted from humans at normal body

temperature• Military use (surveillance, night vision, homing)• Short ranged wireless communication, weather

forecasting, remote temperature sensing– Purple white light get on cheaper digital cameras

(poor infrared filters)

Microwaves

• Wireless LAN & Bluetooth• Radar Detectors, Air Traffic Control• GPS• The frequency of the waves used in microwave

ovens, 2500 megahertz, targets water, sugar & fat molecules– Thin, sharp metals can not handle the electric current

passing through them and may spark or start a fire• Has never been conclusively shown that

microwaves have biological effects• http://www.youtube.com/watch?v=Ug8hSqkFUXY • http://www.youtube.com/watch?v=PIrd4172Czw

Radio Waves

• Transport information through the atmosphere or space without wires

• AM & FM Radio, TV transmission, mobile phones, military communications, wireless computer networks

Visible SpectrumPart of the electromagnetic spectrum

•Continuous Spectrum: One color fades gradually into the next.

•Different colors have different wavelengths.

•The color of visible light with the largest wavelength and lowest frequency is:

•The color of visible light with the shortest wavelength and highest frequency is:

•The brightness of visible light is determined by:

• Radiation with the largest wavelength and lowest frequency is:

• Radiation with the shortest wavelength and highest frequency is:

• Radiation with frequencies greater than visible light can pose health hazards because:

• Radiation with frequencies lower than visible light are less harmful because:

Aircraft & Shipping Bands, Radio Waves

Gamma Rays

Have high enough energy to be capable of damaging organisms

Do not have enough energy and pose no health hazards

What puzzled scientists about electromagnetic radiation?

• Why do objects at different temperatures give off different color light?

• Why do different elements emit different colors when heated?

Planck’s Theory

• Suggested that the energy emitted or absorbed by an object is restricted to “pieces” of particular sizes called quanta.– Substances can emit or absorb only certain

amounts of energy (so only certain wavelengths)

– Showed that frequency and energy are directly proportional

Planck’s Theory

• E = h × – h = Planck’s constant = 6.626 × 10-34 J●s

• Joule (J) = S.I. Unit for Energy

Planck’s Theory

How much energy is contained in a wave with a frequency of 2 x 1016 Hz?What is the frequency of a wave with an energy of 2.90 x 1022 J?

What is the approximate energy of Ultraviolet Light?

What is the energy of radiation with a wavelength of 290 nm?

Why can’t you see quantum effects in the everyday world around you?

• Planck’s constant is very small, therefore, each quantum of energy is very small– Quanta are too small to see in the everyday

world– Atoms are very small, so in relation to the

atom, quanta are significant

The Photoelectric Effect

• Proposed by Albert Einstein – In the photoelectric effect,

electrons are ejected from the surface of a metal when the metal absorbs photons

– Photon: Quantum of light (a tiny particle of light)

The Photoelectric Effect

• When a photon strikes the surface of a metal, it transfers its energy to an electron in a metal atom– If the energy of the photon

is too small for the electron, the electron stays put

– If the photon has enough energy, the electron will escape the surface of the metal.

The Photoelectric Effect

• Why is it easier for violet light (vs. red light) to cause the photoelectric effect?

Violet light has a higher frequency and, therefore, more energy than red light.

Light has a Dual Nature

• A photon behaves like a particle but always travels at the speed of light and has an associated frequency and wavelength– In 1923, Arthur Compton

showed that a photon could collide with an electron

– Light possesses the properties of both particles and waves

How can atoms gain or lose energy?

• Atoms can only gain or lose energy in a quantum

• Take a look through your spectral tube at the emission tube at the front of the room.– How does what you’re looking at demonstrate

the idea above?

Line Spectrum

• A spectrum that contains only certain colors, or wavelengths

•When heat or electricity is passed through an atom, the atom absorbs the energy and then gives off that energy in the form of light

•The emitted light is unique for every element

•Atomic Emission Spectrum: An atomic fingerprint showing the emission line spectrum of that atom

•Useful in identifying an element

NIELS BOHR—1913PLANETARY MODEL OF THE ATOM

Electrons move in defined

orbits around the

nucleus—just as the

planets move

around the sun.

Orbit: Region outsid

e the

nucleus where electro

ns are

found

Bohr’s Postulate

• Was applicable only to hydrogen

• Able to show that electrons move to higher energy levels (excited states) when they absorb radiation.

• Electrons will immediately return to the lower energy levels (ground state) by emitting energy of a specific wavelength

Light has a Dual Nature

– When light travels through space, it acts like a wave

– When light interacts with matter, it acts like a particle

– De Broglie predicted matter waves--that matter should behave like waves and exhibit a wavelength

– Clinton Davisson & Lester Germer proved that electrons (believed to be particles) were reflected from a matter like waves

• Mass of an object must be very small in order to observe its wavelength

Heisenberg Uncertainty Principle

• An electron is located by striking that electron with a photon which bounces back to a detection device

• The electron is so small in mass that the electron is moved by the collision

• Proved a problem with Bohr’s model: You cannot think of electrons moving in defined paths because there is no way to prove the electrons follow defined paths

MODELS OVER TIME

Quantum Mechanical Modela.k.a: Wave Model

•Explains the properties of atoms by treating the electron as a wave that has quantized its energy

•Does not describe exact positions of the electrons; instead describes the probability that electrons will be found in certain locations around the nucleus

Electron CloudAn illustration that uses a blurry cloud to illustrate the probability of finding an electron in various locations

around the nucleus.

(Determined by wave functions electron density charts)

Areas of high

electron density are

the most probable

locations of the

electrons.

Areas of low electron density are the least probable locations of the electrons.

Atomic Orbitals•Region of space where the electron is located

•Have characteristic shapes, sizes and energies

•Do not describe how the electron actually moves

•The orbital occupied is determined by the amount of energy of an electron

s-Orbital

The s-orbital consists of 1 orbital on all 3

axes

1 orbital has a maximum of 2

electrons

p-Orbital

A p-orbital has a dumbbell shape

The p-orbital can exist on 3 different axes (x, y and

z). Therefore there are 3 p orbitals.

The p-sublevel’s 3 orbitals can hold a maximum of 6

electrons (2 on each of the 3 orbitals).

d-Orbital

A d-orbital has a cloverleaf shape

There are 5 different orientations of a d-orbital.

The d-sublevel’s 5 orbitals can hold a maximum of 10 electrons (2 electrons on

each orbital).

f-Orbital

An f-orbital has a complex shape

There are 7 different orientations of the f-orbital.

The f-sublevel can hold a maximum of 14 electrons

(2 for each orbital).

Energy & Orbitals•Energy of electrons are quantized (exact)

•Principal Energy Levels or Principal Quantum Number designates the distance of the electron from the nucleus

Principal energy levels are divided into sublevels

SublevelsSublevels of the

atom are designated:

s, p, d & f

The number of the energy level tells you how many sublevels are present within that sublevel. Another words:

Energy Level 1 has __________ Sublevel

Energy Level 2 has __________ Sublevels

Energy Level 3 has __________ Sublevels

Energy Level 4 has __________ Sublevels

The electrons address consists of its principal energy level, its sublevel, and its electrons within that sublevel

12

34

SUBLEVEL s

Orbital Shape Max # of electrons

Region on Periodic

Table

Orbital Models

1 orbital

s

Sphere 2 Groups 1 & 2

(1st tower)

SUBLEVEL pOrbital Shape Max # of

electronsRegion on Periodic

Table

Orbital Models

3 orbitals

px

py

pz

dumbbell 6 Groups 13-18

(2nd tower)

SUBLEVEL dOrbital Shape Max # of

electronsRegion on Periodic

Table

Orbital Models

5 orbitals

dxy

dxz

dyz

dx2-y2

dz2

cloverleaf 10 Groups 3-12

(transition metals)

SUBLEVEL f

Orbital Shape Max # of electrons

Region on Periodic

Table

Orbital Models

7 orbitals complex 14 Bottom 2 rows (inner-

transition metals)

Some Atomic Models

More Models

Example

Beryllium: ______ protons, ______ electrons

E- Configuration: 1s22s2

Example

Oxygen: ______ protons, ______ electrons

E- Configuration: 1s22s22p4

PRACTICE PROBLEMSElectron configurations for 3 different elements are given below. Draw the atomic model of each element and then identify the element.

Examples: 1s22s1 1s22s22p3 1s22s22p63s23p4

1) 1s22s22p1 2) 1s2 3) 1s22s22p63s1

ExampleBoron: ______ protons, ______ electrons

E- Configuration: 1s22s22p1

Examples

Helium: ______ protons, ______ electrons

E- Configuration: 1s2

Examples

Sodium: ______ protons, ______ electrons

E- Configuration: 1s22s22p63s1

Electron Spin

•Electrons spin either clockwise or counterclockwise

•The spinning creates a magnetic field

•Clockwise is like a magnet whose north pole is pointing up

•Counterclockwise behaves like a magnet whose north pole is pointing down

•Parallel Spins result in a net magnetic effect

•Opposite Spins cancel each other out

Pauli Exclusion Principle

-1925, Austrian physicist-Wolfgang Pauli

-States that each orbital in an atom can hold at most 2 electrons and that these electrons must have opposite spins (or be paired).Sublevels Orbitals Max # of e-

s 1 2

p 3 6

d 5 10

f 7 14

Electron Configuration• The addresses of an atom’s electrons• Determined by distributing the atom’s electrons

among levels, sublevels and orbitals based on a set of principles

• Orbitals from lowest to highest energy:

s p d f• Ground State: The electrons are in the lowest

energy levels available

How do electrons occupy energy levels?

• Aufbau Principle: Electrons are added one at a time to the lowest energy orbitals available until all the electrons are accounted for

• Pauli Exclusion Principle: An orbital can hold a maximum of 2 electrons that must spin in opposite directions

• Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results

Orbital Diagrams

4p ____ ____ ____3d ____ ____ ____ ____ ____4s ____3p ____ ____ ____3s ____2p ____ ____ ____2s ____1s ____

What happens when an element in its ground state is supplied with electricity

or heat?– Electrons may move to the excited state.– Excited State: Energy level attained when an electron

absorbs energy and jumps to a higher energy level

Ground State

Excited State

For each pair of orbital diagrams below, which represents the ground state and which

represents the excited state of that atom?4p ____ ____ ____3d ____ ____ ____ ____ ____4s ____3p ____ ____ ____3s ____2p ____ ____ ____2s ____1s ____

Scandium

Magnesium

4p ____ ____ ____3d ____ ____ ____ ____ ____4s ____3p ____ ____ ____3s ____2p ____ ____ ____2s ____1s ____

4p ____ ____ ____3d ____ ____ ____ ____ ____4s ____3p ____ ____ ____3s ____2p ____ ____ ____2s ____1s ____

4p ____ ____ ____3d ____ ____ ____ ____ ____4s ____3p ____ ____ ____3s ____2p ____ ____ ____2s ____1s ____

What happens to the excited electron?

http://www.meta-synthesis.com/webbook/11_five/five04.jpg

Exceptions to the Aufbau Rule

• A half-full or full d sublevel will increase an atom’s stability– An electron may be removed from the s

sublevel to create a full or half full d sublevel4p ____ ____ ____3d ____ ____ ____ ____ ____4s ____3p ____ ____ ____3s ____2p ____ ____ ____2s ____1s ____

4p ____ ____ ____3d ____ ____ ____ ____ ____4s ____3p ____ ____ ____3s ____2p ____ ____ ____2s ____1s ____

Cr Cu

Groups (also called Families)

• The vertical columns on the periodic table

• There are 18 groups, labeled with the numbers 1-18.

1

2

3 4

15

5 76 98

17 16

18

1413

1211 10

d-block Transition MetalsGroups 3-12Begins with Principal Energy Level 3Contains 10 elements

because each d sublevel can hold up to 10 electrons

Corresponding Regions on the Periodic Tables-block

Representative Elements

Groups 1 & 2

Begins with Principal

Energy Level 1

Contains 2 elements

because each s sublevel

can hold 2 electrons

p-block Representative ElementsGroups 13-18Begins with Principal Energy Level 2Contains 6 elements because each p sublevel

can hold up to 6 electrons

f-block

Inner Transition Metals

lanthanides & actinides

(bottom 2 rows)

Begins with Principal

Energy Level 4

Contains 14 elements

because each f sublevel

can hold up to 14

electrons

He

N

Ti

I

Ce

Fr

Noble Gas Configuration: Uses the symbol of the noble gas in brackets to represent the inner level electrons of an atom.

Cd

U

Ba

1s

3s

2s

6d

5d

4d

3d

5f

4f

7p

6p

5p

4p

3p

2p

4s

5s

6s

7s

VALENCE ELECTRONS

The electrons in the outermost energy level.

Remember, the number in front of the sublevel indicates the energy level:

1s22s22p6

So…find the highest energy level and add up all the electrons in that level.

EXAMPLES

• Calcium

• Aluminum

• Iodine

• Oxygen

• Iron

ENERGY

Electrons with the most energy are located farthest from the nucleus

Electrons with the lowest

energies are located close

to the nucleus.

Quantum

A quantum is the specific amount of energy needed for

an electron to move between energy

levels.