Electrons Galore!

Electrons and Light• By 1900, scientists knew that light could be thought of as moving

waves that have given frequencies, speeds, and wavelengths.• In empty space, light waves travel at 2.998 × 108 m/s. • The wavelength is the distance between two consecutive peaks or

troughs of a wave. – The distance of a wavelength is usually measured in meters. – The wavelength of light can vary from 105 m to less than 10−13 m.– This broad range of wavelengths makes up the electromagnetic spectrum

• Our eyes are sensitive to only a small portion of the electromagnetic spectrum. – This sensitivity ranges from 700 nm, which is about the value of

wavelengths of red light, to 400 nm, which is about the value of wavelengths of violet light.

Electromagnetic Spectrum

Photoelectric Effect• In 1905, Albert Einstein proposed that light also has some properties of

particles. • His theory would explain a phenomenon known as the photoelectric effect.

– This effect happens when light strikes a metal and electrons are released.– What confused scientists was the observation that for a given metal, no electrons

were emitted if the light’s frequency was below a certain value, no matter how long the light was on.

– Yet if light were just a wave, then any frequency eventually should supply enough energy to remove an electron from the metal.

• Einstein proposed that light has the properties of both waves and particles.• According to Einstein, light can be described as a stream of particles, the

energy of which is determined by the light’s frequency. • To remove an electron, a particle of light has to have at least a minimum

energy and therefore a minimum frequency.

Light is a wave

• When passed through a glass prism, sunlight produces the visible spectrum– all of the colors of light that the human eye can see. You can

see• Remember that the visible spectrum is only a tiny

portion of the electromagnetic spectrum. • The electromagnetic spectrum also includes X rays,

ultraviolet and infrared light, microwaves, and radio waves.

• Each of these electromagnetic waves is referred to as light, although we cannot see these wavelengths.

Frequency, wavelength and Energy

• Red light has a low frequency and a long wavelength.

• Violet light has a high frequency and a short wavelength.

• The frequency and wavelength of a wave are inversely related.

I see the light!

• When a high-voltage current is passed through a tube of hydrogen gas at low pressure, lavender-colored light is seen. When this light passes through a prism, you can see that the light is made of only a few colors.– This spectrum of a few colors is called a line-emission spectrum.

• Experiments with other gaseous elements show that each element has a line-emission spectrum that is made of a different pattern of colors.

• In 1913, Bohr showed that hydrogen’s line-emission spectrum could be explained by assuming that the hydrogen atom’s electron can be in any one of a number of distinct energy levels. – The electron can move from a low energy level to a high energy level by

absorbing energy.

I still see the light!

• Electrons at a higher energy level are unstable and can move to a lower energy level by releasing energy. – This energy is released as light that has a specific wavelength. – Each different move from a particular energy level to a lower energy

level will release light of a different wavelength.

• Bohr developed an equation to calculate all of the possible energies of the electron in a hydrogen atom. – His values agreed with those calculated from the wavelengths

observed in hydrogen’s line-emission spectrum. – In fact, his values matched with the experimental values so well that

his atomic model that is described earlier was quickly accepted.

Light tells me where electrons live!

• Normally, if an electron is in a state of lowest possible energy, it is in a ground state

• If an electron gains energy, it moves to an excited state

• An electron in an excited state will release a specific quantity of energy as it quickly “falls” back to its ground state.

• This energy is emitted as certain wavelengths of light, which give each element a unique line-emission spectrum

Electron Transitions

Quantum Numbers

• The present-day model of the atom, in which electrons are located in orbitals, is also known as the quantum model.

• According to this model, electrons within an energy level are located in orbitals, regions of high probability for finding a particular electron.

• However, the model does not explain how the electrons move about the nucleus to create these regions.

• To define the region in which electrons can be found, scientists have assigned four to each electron.

n tells you in which period the atom is

• The principal quantum number, symbolized by n, indicates the main energy level occupied by the electron.

• Values of n are positive integers, such as 1, 2, 3, and 4.

• As n increases, the electron’s distance from the nucleus and the electron’s energy increases.

l tells you the shape of the orbital

• The main energy levels can be divided into sublevels. • These sublevels are represented by the angular momentum

quantum number, l. • This quantum number indicates the shape or type of orbital that

corresponds to a particular sublevel. There are letter codes for this quantum number– l = 0 corresponds to an s orbital– l = 1 to a p orbital – l = 2 to a d orbital– l = 3 to an f orbital.

• For example, an orbital with n = 3 and l = 1 is called a 3p orbital, and an electron occupying that orbital is called a 3p electron.

s Orbitals (l = 0)—Shape

• Each shell has one s orbital.• Lowest energy orbital in a shell.• Spherical in shape.

s Orbitals (l = 0)—Shape • For n = 2 and l = 0

– 2s orbital• For n = 3 and l = 0

– 3s orbital• The larger the value of n, the

larger the orbital.– The further out from the nucleus

it extends.• Contains nodes-regions of no

probability.– number of nodes = (n – 1)

p Orbitals (l = 1)—Shape • Each principal energy state n = 2 and above has three p orbitals.

– ml = -1, 0, +1• Each of the 3 orbitals point along a different axis.

– px, py, pz

• 2nd lowest energy orbitals in a principal energy state.• Shape has two lobes.• Node at the nucleus. Total of n nodes.

d Orbitals (l = 2)—Shape • Each principal energy state n = 3

and above has five d orbitals– ml = -2, -1, 0, +1, +2

• 4 of the 5 orbitals are aligned in a different plane– the fifth is aligned with the z axis,

dz squared

– dxy, dyz, dxz, dx squared – y squared

• 3rd lowest energy orbitals in a principal energy state

• Mainly 4-lobed.– one is two-lobed with a toroid

f Orbitals (l = 3)—Shape • Each principal energy state

n = 4 and above has seven f orbitals– ml = -3, -2, -1, 0, +1, +2, +3

• 4th lowest energy orbitals in a principal energy state.

• Mainly 8-lobed– some 2-lobed with a toroid

mL tells you the orientation

• The magnetic quantum number, symbolized by m, is a subset of the l quantum number.

• It also indicates the numbers and orientations of orbitals around the nucleus.

• The value of m takes whole-number values, depending on the value of l.

• The number of orbitals includes one s orbital, three p orbitals, five d orbitals, and seven f orbitals.

ms tells you the magnetic nature

• The spin quantum number is symbolized by ±½ • (↑ or ↓), indicates the orientation of an

electron’s magnetic field relative to an outside magnetic field.

• A single orbital can hold a maximum of two electrons, which must have opposite spins.

Electron Configurations• Each orbital can hold a maximum of two electrons. • The discovery that two, but no more than two, electrons can occupy a

single orbital was made in 1925 by the German chemist Wolfgang Pauli. – This rule is known as the Pauli Exclusion Principle

• Another way of stating the Pauli exclusion principle is that no two electrons in the same atom can have the same four quantum numbers.

• The two electrons can have the same value of n, the same value of l and they may

• have the same value of m. But these two electrons cannot have the same spin quantum number. – If one electron has the value of +½ then the other electron must have the

value of -½

Electron Configurations

• The arrangement of electrons in an atom is usually shown by writing an electron configuration.

• Like all systems in nature, electrons in atoms tend to assume arrangements that have the lowest possible energies.

• An electron configuration of an atom shows the lowest-energy arrangement of the electrons for the element.

Electron Configuration of Atoms

• Every electron in an atom has a unique set of quantum numbers.– Quantum numbers describe which orbitals are

occupied by electrons in an atom.• Only two electrons can occupy one orbital.

• Arrangement of electrons an atom is called the electron configuration.– lowest energy arrangement = ground state

Electron Configuration of Atoms

• Two ways to depict the electronic structure of an atom.

• Orbital filling diagram.– Follow Aufbau Principle.

• Electron configuration.– Full electron configuration.– Condensed using the “inner electron

configuration” based on the closest previous Noble Gas.

Rules of the Aufbau Principle• Determine number of electrons in atom or ion.• Orbitals are filled in order of increasing energy, lower energy orbitals first.

– Lowest energy shells fill first.• n =1 before n = 2, etc.

– Lowest energy orbitals in a shell fill first.• s p d f

• An orbital can hold two electrons of opposite spin—Pauli Exclusion Principle.

• For degenerate orbitals (same energy), the lowest energy is attained when the number of electrons with the same spin is maximized—Hund’s Rule.– Half fill orbitals of the same energy before pairing electrons.

• One electron per orbital until subshell is half full.– All have same spin, prior to pairing.

Electron Configuration Example

1s

2s

2p

3s

3p3d4s

4p4d

5s

5p

Na Atomic Number (Z): 11

1s2 2s2 2p6 3s1 or [Ne] 3s1

n = 1 l = 0

n = 2 l = 0

n = 2 l = 1

n = 3 l = 0

1s 2p2s 3s

n = 3 l = 1

Electron Configuration MoreIdentify the element with the ground state configuration

1s2 2s2 2p4

A. Carbon

B. Oxygen

C. Sulfur

D. Argon

Electron Configuration Your Turn

• Draw the ground state orbital filling diagram for vanadium.

A. [Ar] _ _ _

B. [Ar] _ _ _ _ _ __ .

C. [Ar] _ _ _ _ _

D. [Ar] _ _ _ __ __

4s 4p

4s 4p

4s 3d

3d

3d

Electron Configuration and the Periodic Table

s block p block d block f block

Beginhere

Endhere

Beginhere

Aufbau Principle

Endhere

2p - 2p3p - 3p4p - 4p5p - 5p6p - 6p7p - 7p

3d - - - 3d4d - - - 4d5d - - - 5d6d - - - 6d

4f - - - - - 4f5f - - - - - 5f

1s2s3s4s5s6s7s

Block of elements by last filled atomic orbitals

2p - 2p3p - 3p4p - 4p5p - 5p6p - 6p7p - 7p

3d - - - 3d4d - - - 4d5d - - - 5d6d - - - 6d

4f - - - - - 4f5f - - - - - 5f

1s2s3s4s5s6s7s

Block of elements by last filled atomic orbitals

Electron Configuration andthe Periodic Table

• The group number corresponds to the number of valence electrons.

• The Period number corresponds to the principle energy level (n) of the valence electrons.

• The length of each “block” is the maximum number of electrons the subshell can hold.– Subshell = same orbital type in a level

• Example: all p orbitals (l = 1) in n = 3 level.

More Electron Configuration

• What is the ground state electron configuration for germanium (Ge)?

A. [Ar] 3d10 4s2 4p2

B. [Ar] 4s2 4p6 3d6 C. [Ar] 4s2 3d10 4p2

D. [Ar] 3d10 4s2 4p1

More Electron Configuration

• What is the ground state electron configuration for tantalum (Ta)?

A. [Xe] 6s2 5d3

B. [Xe] 6s2 4f14 5d3 C. [Xe] 6s2 5d10 4p7

D. [Xe] 5d10 6s2 4p7