1

What have we learned ?

1) Atoms are the building blocks of matter.2) Atoms are composed of a dense nucleus

ELECTRONIC STRUCTURE OF ATOMS/ QUANTUM THEORY (CHAPTER 7)

) m mp f(protons and neutrons) with electrons far removed from the nucleus.

3) Stoichiometry.4) Elements are arranged in a periodic table.5) Behavior of gases (laws and theory). 6) Basics of thermochemistry.

What we don't know ?

1) Why atoms combine to form molecules ?2) Why the periodic table has its particular arrangement of the elements ?3) What is the detailed structure of an atom 3) What is the detailed structure of an atom (electrons) ?4) Why atoms combine to form compoundswith particular formulas ?5) Why elements (main group, Group A) form stable monatomic ions with a particular charge (e.g., Na+, Ca2+, O2-, Cl-) ?

MOLECULAR FORMULAS

• Atoms combine to form molecules with specific formulas.

• How atoms combine depends on the distribution or arrangement of electronsaround the nucleus of the atom (i.e., l i )electronic structure).

• Early experimental evidence aided in developing an understanding of atomic structure:– J. J. Thomson’s discovery of the electron.– Rutherford’s Nuclear Model of the Atom.

RELATIONSHIP BETWEEN ATOMS AND ENERGY

A GOOD EXAMPLE IS:

Atoms + energy emit light

Different elements (atoms) emit light of different colors (red, green, yellow, orange, blue).

M* M + lightexcited emissionatom

M + light M* absorption excited

Light interacts with atoms in specific ways –gain information on the electronic structure of atoms by studying this interaction.

absorption excited atom

Light is electromagnetic radiation and is a wave and it has wavelike characteristics.

Wave: Vibrating disturbance by which energy

Waves are characterized by: frequencywavelengthamplitude

Wave: Vibrating disturbance by which energyis transmitted.

Water waves, sound waves, etc.

2

Figure 7.2

WAVE CHARACTERISTICS

Wavelength ( lambda): The distance(length) between adjacent peaks on the wave.SI unit: meter (m).

Frequency ( nu): The number of waves q y(wavelengths) that pass a point per second.The SI unit of frequency is the hertz (Hz).SI units: 1/s or s-1 or hertz (Hz).

Amplitude (intensity): The vertical distancefrom the mid-line to the peak (intensity).

Figure 7.2

Frequency/Amplitude(intensity) of Various Waves

Speed of wave propagation = X

speedtime

1distance

time

distance speed

Electromagnetic radiation (radiant energy, light).

Speed of light (c) = 3.00 X 108 m/s (in a vacuum)p g ( ) ( )(186,000 miles/s or 671,000,000 mph)

This radiation does not need a medium for propagation - can travel through a vacuum.

c = = constant

If increases, then decreases (inverse relationship).

3

Example: Frequency and Wavelength

What is the frequency () of radiation that has a wavelength () of 1.5 X 10-2 m ?

Start this problem with c =

R h i d l f Rearrange the equation and solve for :

Hz 10 X2or

s 10 X2 m 10 X1.5

m/s 10 X3.00

10

1-102-

8

c

Electromagnetic Radiation

Maxwell (1873): Developed a theory for light.

He proposed that light waves consists of two components; a vibrating electric field and a vibrating magnetic field.vibrating magnetic field.

Hence, the term “electromagnetic radiation”.

This radiation involves the transmission of energy from one place to another.

Figure 7.3

Figure 7.4

Max Planck (1900)

Studied radiation emitted by matter (objects). All objects emit electromagnetic radiation.

Postulated that atoms and molecules can emit or absorb energy (radiation) only in discrete quantities. Energy in matter is quantized, not continuous !!

4

QUANTIZED ENERGY

What is a quantum of energy ?– The smallest increment of energy that can be emitted or absorbed in the form of electromagnetic energy (light) by matter.

Energy (E) of a single quantum is given by:E = h

E = energy (J) = frequency (1/s)h = Planck's constant (6.63 X 10-34 J · s)

– Units: E = (J · s)(1/s) = J (energy)

PHOTOELECTRIC EFFECT: Light can cause electrons to be ejected from a metal surface.

lightelect ron

met al surface

There is a threshold frequency () for electron ejection and if light has a less than the threshold, no electrons are ejected. At a greater than the threshold, the number of electrons ejected is proportional to the light intensity.

PHOTOELECTRIC EFFECT

These observations of the photoelectric effect could not be explained by the wave theory of light.

Alb t Ei t i (1905) C id d th t li ht i Albert Einstein (1905): Considered that light is composed of a stream of particles that are called photons (or quanta-packets of energy).

Extended Planck’s idea’s.

Photoelectric Effect

Figure 7.5

5

Each photon of light has an energy, E:

E = h · (Planck’s ideas)Energy of a photon is directly proportional to the frequency.

Energy of a photon is important and dictates gy f p phow light interacts with matter.

Electromagnetic radiation (light) has both wave-like and particle-like properties.

The dual nature of light!!!

Problem: Quantized EnergyA typical laser pointer has red light -wavelength 630 nanometers (6.3 X 10-7 meters).

What is the energy of one photon of light with = 630 nm?

J 10 X3.16

m) 10 X(6.3

m/s) 10 Xs)(3.0J 10 X(6.63

λν

19-

7-

834-

hc

h E

Electromagnetic SpectrumRadio waves

Microwave

InfraredInverse relationship between and

Some Facts

Visible

Ultraviolet

X-rays

Gamma rays

(lambda) ranges from 1013 to 10-3 nm

(nu) ranges from 104 Hertz (Hz) to 1020 Hz

Figure 7.4

Energy of Electromagnetic Radiation

• Given by E = h = hc/• Short wavelengths () high energy; e X rays very dama in to life e.g., X-rays - very damaging to life

• Long wavelengths () low energy; e.g., radio waves - safe

ATOMIC SPECTRAExperimental observations of the interaction of radiation with matter.

Atom + h Atom* absorptionAtom* Atom + h emission

Atoms absorb and emit light with discrete energy(’s and ’s). They can have only particular amounts of energy – energy of atoms is quantized.

Energy levels of atoms - it's the electrons that are important - energy levels of the electrons.

6

Atomic Spectrum of Hydrogen Atoms

Simplest element, it has the simplest spectrum.

See a series of lines in the atomic spectrum of hydrogen atoms.

O l di t ’ d ’ f li ht b b dOnly discrete ’s and ’s of light are absorbedor emitted by hydrogen atoms.

This resulting spectrum is unique for hydrogen -different from all other elements.

Experiment to View the Atomic Spectrum of H

Figure 7.6

H2 H H* H + lightenergy energy emission

Continuous Spectrum

Each type of atom (element) has its own unique spectrum (lines).

Figure 7.6

Atomic Spectra of Hydrogen

Numerous lines ranging from the ultraviolet ( < 400 nm) to the infrared region ( > 800 nm).

J. J. Balmer (1885): Studied the lines in the visible portion of the the spectrum (400 nm to 700 nm).

F d i l ti ( i i l) f l l ti

1

2

1

λ

122H

nR RH = Rydberg constant

= 109,678 cm-1

n = 3, 4, 5, and 6

Found a simple equation (empirical) for calculating the emission wavelengths in the visible region:

7

Atomic Spectra of Hydrogen

General equation for calculating all emissionlines in the hydrogen spectrum:

RH = Rydberg constant= 109,678 cm-1

22H11

λ

1R ,

Restriction: ni and nf are whole numbers ranging from 1 to infinity.

The above equation is called the Rydberg equation.

22λ fi nn

Different Series of Spectral Lines for H atoms

Lyman series: nf = 1; ni = 2, 3, 4,…(ultraviolet)Balmer series: nf = 2; ni = 3, 4, 5,…(visible)Paschen series: nf = 3; ni = 4, 5, 6, …(infrared)Brackett series:nf = 4; ni = 5, 6, 7, …(infrared)

EXPLANATION OF H-LINE SPECTRUM

Niels Bohr (1913), Danish physicistDeveloped a theoretical model for the hydrogen atom.

Bohr viewed the electron as moving around the Bohr viewed the electron as moving around the nucleus in only fixed (discrete) radii or orbits (planets in a solar system).

This placed restrictions on the sizes of the orbitsand the energy that the electron can have in a given orbit (departure from classical physics).

Will the photon be absorbed ?

1st Bohr orbit

2nd Bohr orbit

3rd Bohr orbit

If of the proper energy, the photon is absorbed and the electron jumps to a higher energy orbit.

Absorption

Electron can return to the ground state by emitting a photon of the same energy.

1st Bohr orbit

2nd Bohr orbit

3rd Bohr orbit

Emission

1st Bohr orbit

2nd Bohr orbit

3rd Bohr orbit

8

Figure 7.9

Bohr described the potential energy of the electron in the hydrogen atom by a simple (?) equation:

m = mass of the electron (9 109 X 10-31 kg)

22

422

hn

meπ E

m = mass of the electron (9.109 X 10 kg)e = charge of the electron (1.602 X 10-19 C,

(1.519 X 10-14 kg1/2 m3/2 s-1)h = Planck’s constant (6.63 X 10-34 J · s)n = quantum number (1, 2, 3, ….)

J 10 X2.182 -18

2

42

H h

meπ R'

Bohr’s Equation of Energy

R’H: constant (2.18 X 10-18 J)n: quantum number

(has values from 1 to )

(integers) ... 3, 2, 1,2H

n

n

R'-E n

(has values from 1 to )

Negative sign: Potential energy of the atom < potential energy of H+ and a free electron.

Lowest energy, n = 1 (E < 0)Highest energy, n = (E = 0)

Potential Energy

Bohr’s Model of H-atomn = 1: First Bohr orbit, ground state, lowest potential energy.

If the hydrogen atom in the ground state absorbs energy, then the electron goes to a higher energy (larger) orbit (n = 2, 3, 4, ...); g gy ( g ) ( )called excited states.

For transitions in the Bohr Model of the H-atom

n = 1 to n = 2 absorbs energy

n = 2 to n = 1 emits energysame energy

Bohr’s Hydrogen Atom

b R’H

(integers) ... 3, 2, 1,2H

n

n

R'-E n

9

Figure 7.9

Energy of Transitions in Bohr H-Atom

Energy of the Initial State:

2

H

ii n

R'E

Energy of Final State:

2H

ff n

R'E

hνR'E

n

R'-

n

R'-EEE

ifif

2H

2H

11Δ

Δ

Energy of Transitions in Bohr H-Atom

If ni > nf , then a photon is emitted, E is negative.If ni < nf , then a photon is absorbed, E is positive.

hνnn

REfi

22HΔ

Example: What is the change in energy of the hydrogen atom when the electron goes from the n = 3 to the n = 1 level ?

Use the equation for the change in energy for H-atom:

11Δ H

R'E

A photon of light energy is emitted, since ni

> nf , E is negative.

J 10 X1.941

1

3

1J 10 X2.18Δ

Δ

-1822

-18

22H

E

nnRE

fi

What is the frequency () and wavelength () of this photon ?

Use Einstein‘s equation and solve for :

Hz 10 X2.93sJ10X6 63

J 10 X1.94Δν

νΔ

1534-

18-

h

E

hE

Calculate wavelength () :

sJ 10X6.63 34 h

nm 102m 10 X1.02 /s10 X2.93

m/s 10 X3.00

νλ

νλ

7-15

8

c

c

10

Relationship between Bohr’s equation and the Rydberg equation:

Bohr’s equation:

22H

11Δ

fi nnR'E

Rydberg Equation:

fi nn

22H11

λ

1

fi nnR

1

1 J incm in

hc

R'R H

H

Rydberg constant calculated from Bohr’s theory:

λ

1

λΔ hc

hc hv E

-1

834-

-18

cm 109,600

cm 100

m 1

m/s 10 X3.00sJ 10 X6.63

J 10 X2.18

Agrees well with the known Rydberg constant value (109,678 cm-1).

Bohr’s Theory• Only works for the hydrogen atom; not a

general theory, it is limited and incomplete.

• Used some elements of "classical physics", which do not apply to atomic and subatomic particles !!!!particles !!!!

• Questions for Bohr:– Why is energy of the electron quantized?– Why are electrons restricted to only certain orbits?

– Why are atoms stable (exist)?

de Broglie’s Hypothesis (1924)• If electromagnetic radiation (light) can have

particle-like properties (photon), then electronsmay also possess wave-like properties.

• Electrons do indeed have both particle and wave-like properties !!! Electrons are similar to light.g

• From Einstein's theory of relativity:

This equation relates energy and mass.light of speed

mass

2

c

m

mcE

de Broglie’s HypothesisUsing photon equation for light:

Substitute mc2 for E and solve for :

c

c

hchE

Substitute mc2 for E and solve for :

mc

h

hcmc

and

2

de Broglie’s Hypothesis

• Notice that the product, mc, is defined as momentum (i.e., mass X velocity).

• This equation then relates wavelength, mass, and velocity.mass, and velocity.

• For an electron, its speed must be less than the speed of light, c. We can simply substitute in for the speed (velocity) of the electron, v.

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The wavelength of an electron (a particle) is given by:

property) (particle momentumv

(m/s) velocityvv

m

m

h

)wavelength (distance, m

smkg

smkg

:Units2

Electrons, protons, and neutrons experience diffraction - wave property !!!!

p p y)(p

Diffraction Pattern for Al with X-rays and Electrons

X-rayselectronsFigure 7.14

Example:Calculate the de Broglie wavelength for an electron travelling at 1 % the speed of light (3.00 X 106 m/s).

Use de Broglie’s relationship: v

m

h

For an electron,

m = 9.11 X 10-31 kg; v = 3.00 X 106 m/s

Follow the units through!!!!

pm 243or m 10 X2.43 m/s 10 X3.00kg10 X9.11

sJ10 X6.63 λ 10-

631-

34-

For the baseball:

de Broglie Wavelength For a Baseball

v

m

h

mass is 0.14 kg and velocity is 40 m/s

m 10x 1.2 m/s) kg)(40 (0.14

sJ10x 6.63 34-

34-

Wave properties of particles only apply to atomic and subatomic particles – not to baseballs !!!!

QUANTUM MECHANICS (WAVE MECHANCIS)

Applies wave properties of matter to explain

Erwin Schrödinger (1926)Nobel Prize (1933)

pp p p patomic properties.

Bohr’s theory: Electrons move around the nucleus in fixed (discrete) orbits (e.g., solar system).

Wave mechanics: Treats electrons as waves –difficult to perceive.

The electron wave function is called an orbital (atomic orbital).

Th t “ bit l” i d t di ti i h f

This theory leads to a series of mathematical functions called wave functions.

The term “orbital” is used to distinguish from Bohr’s ideas of fixed orbits.

12

Interference of Waves

Two waves with the same frequency add together if they are “in phase” (their peaks exactly match).

Constructive interference.

Interference of Waves

Two waves with the same frequency but exactly “out of phase” (peak of one wave matches a valley f h h ) l of the other wave) cancel

out (destroyed).

Destructive interference.

4 wavelengths –constructive interference

Allowed (Stable) Orbital:

4.5 wavelengths –destructive interference

Forbidden (Unstable) Orbital

Electron Waves

4 wavelengths 3.6 wavelengths

destructiveinterference

These electron wave functions are quantized. They must have integral numbers of wavelengths, otherwise destructive interferencewill result in annihilation of the electron wave.

Energy changes in an atom - simply changes in wave functions (patterns) which are quantized.

Electrons exist only in those regions in which their waves reinforce each other.

Rather, the electron wave function takes on a particular shape, a probability of finding the electron in a particular volume of space.

In wave mechanics, the electrons do not orbit the nucleus in fixed, circular orbits (Bohr’s ideas).

An atomic orbital is characterized by:energysizeshapeorientation in space

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TYPES (SHAPES) OF ATOMIC ORBITALSThe “s” Orbitals: Have a spherical shape:

The “s” Orbitals: Have a spherical shape:

Figure 7.18

The “s” Orbitals: Have a spherical shape:

The “p” Orbitals: Are not spherical, consist of two lobes of electron density about the nucleus.

There are three individual “p” orbitals, differ only in their orientation in space:

Figure 7.20

The three individual “p” orbitals, represented in an atom:

14

The “d” Orbitals: Have a more complex shape (four lobes of electron density about the nucleus).Five individual “d” orbitals, differ only in their orientation in space:

Figure 7.21

The “f” Orbitals: Have a very complex shape (eight lobes of electron density about the nucleus).There are seven individual “f” orbitals that varyonly in their orientation in space.

Quantum Numbers

Bohr’s theory used a single quantum number (n) to describe an orbit.

Quantum mechanics theory (wave mechanics) has three quantum numbers (n, l, ml) to describe an atomic orbital.

QUANTUM NUMBERS

• PRINCIPAL QUANTUM NUMBER (n) Determines the energy and size of an orbital.– Allowed values 1, 2, 3,… (positive integers)

• As n increases the energy and size of the As n increases, the energy and size of the orbital increases.

• All orbitals that have the same value for "n" are in the same "shell".

n = 1 shell (first shell)n = 2 shell (second shell)n = 3 shell (third shell)

etc.

Angular Momentum (Shape) Quantum Number (l): This quantum number determines the shape of the orbital.

Divides a shell into smaller groups called subshells.Only certain "l" values are possible-limited by "n":Can have integral values from zero to n – 1

(0 1 2 3 1)(0, 1, 2, 3, ....n - 1)

Shells Subshellsn = 1 l = 0

n = 2 l = 0, 1

n = 3 l = 0, 1, 2

Angular Momentum (Shape) Quantum Number (l)Typically, a letter code is used to specify the “l” values:

l 0 1 2 3 4 5….letter s p d f g h….

Letter designations Nature of spectral feature:

s Sharp-very narrowp Principle-very strongd Diffuse-spread outf Fundamental

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To designate a subshell, we need to specify both "n" (number) and "l" (letter):

Shell Subshell designation

Angular Momentum (Shape) Quantum Number (l)

n = 1 l = 0 1s

n = 2 l = 0 2s

n = 2 l = 1 2p

n = 3 l = 2 3d

• For a given "n" value, there are small differences in energy for various "l" values.

Angular Momentum (Shape) Quantum Number (l)

• Order of increasing energy of subshells in the same shell:

s < p < d < f < g….

4s < 4p < 4d < 4fincreasing energy

MAGNETIC QUANTUM NUMBER (ml)

• Describes the orientation in space relative to other orbitals of that same subshell (l). – Divides a subshell into individual orbitals.

• Possible ml values :(-l, -l + 1, ....+l)• For the "s" subshell:

l = 0 then = 0l = 0 then ml = 0• There is only a single orbital in an “s” subshell.• For the "p" subshell:

l = 1 then ml = -1, 0, 1(three ml values)

• There are three orbitals in a "p" subshell.

MAGNETIC QUANTUM NUMBER (ml):

• For a "d" subshell:l = 2 then ml = -2, -1, 0, 1, 2

(five ml values)

Th five bit l i "d" b h ll• There are five orbitals in a "d" subshell.

• For the "f" subshell:l = 3 then ml = -3, -2, -1, 0, 1, 2, 3

(seven ml values)

• There are seven orbitals in a "f" subshell.

REPRESENTATION OF ORBITALS (shapes and sizes)

Electron wave form (function) considers a probability of finding the electron around the nucleus.

Define the size and shape of an orbital by a 90% probability diagram (90% of the electron density lies within these boundaries, 90 % chance that the electron will be found within this volume).

16

The “s” Orbitals: n = 1 to infinity; l = 0; ml = 0

Have a spherical shape:

1s < 2s < 3s

Increasing size and energy

The “p” Orbitals:

n 2; l = 1; ml = -1, 0, 1

Three individual orbitals in a “p” subshell.

They all have the same shape, size, and energy differ only in their orientation in energy - differ only in their orientation in space.

Consist of two lobes of electron density about the nucleus.

There are three individual “p” orbitals in a “p”subshell, differ only in their orientation in space:

The “d” Orbitals:

n 3; l = 2; ml = -2, -1, 0, 1, 2

Five individual orbitals in a “d” subshell.

They have a more complex shape, differ only in their orientation in space.

Each have the same energy, size, and shape.

Five individual “d” orbitals, differ only in their orientation in space:

Figure 7.21

17

THE “f” ORBITALS:

n 4 ; l = 3; ml = -3,-2, -1, 0, 1, 2, 3

Seven individual orbitals in the “f” subshell.Very complex shape (8 lobes of electron density) and differ only in their orientation in space.

Each have the same energy, size and shape.

Orbitals with the same energy, i.e., the three p orbs, five d orbs, etc. are said to be degenerate.

MAGNETIC PROPERTIES OF THE ELECTRON

It was found that atoms of some elements have magnetic fields.

This was explained by considering that the electron is spinning. electron is spinning.

From electromagnetic theory, a spinning chargegenerates a magnetic field.

Hence, electrons behave like very small magnets.

Quantum mechanics theory predicts that the electron has only two possible spin states.

Hence, electron spin is quantized.

A l i b ( ) i

ELECTRON SPIN

An electron spin quantum number (ms) is assigned to the electron.

ms has values of +1/2 and -1/2

These two spin states generate magnetic fieldsthat are directly opposed.

MAGNETIC PROPERTIES OF THE ELECTRON

ELECTRON SPIN

The quantized spin of the electron is verified by experiment.

Figure 7.17

A complete set of quantum numbers to describe an electron’s wave form (function):

n l ml ms

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Pauli exclusion principle: States that no two electrons in an atom can have the same set of four quantum numbers.

n, l, ml specify an orbital

ms specify electron spin

Consequently, no more than two electrons can be in an orbital. If an orbital contains two electrons, then they must have opposite spin quantum numbers (ms values).

n

value

l

value

designation (subshell)

ml

values

1 0 1s 0

2 0 2s 0

Quantum Numbers

2 1 2p -1, 0, 1

3 0 3s 0

3 1 3p -1, 0, 1

3 2 3d -2, -1, 0, 1, 2

Subshell

# of orbitals

maximum # of electrons

s 1 2

p 3 6

Electrons in Subshells

p

d 5 10

f 7 14

Shell (n) Subshells

maximum # of electrons

1 s 2

Electrons in Shells

1 s 2

2 s, p 8

3 s, p, d 18

4 s, p, d, f 32

Quantum (Wave) MechanicsQuantum Numbers:Principal - n (energy & size of orbitals) Angular Momentum - l (shape of orbitals)Magnetic - ml (orientation of orbitals)Spin - ms (spin of e-’s in orbitals)

P li E l i P i i lPauli Exclusion Principle:– No two electrons in an atom can have the same set of four quantum numbers

– n, l, ml specify an orbital – e- spin distinguishes the two e-’s in an orbital (ms)

Electron Configuration (Electronic Structure)

The distribution of electrons among orbitals of an atom.

Ground State electron configuration: Lowest energy arrangement of electrons.

Electron configuration: List subshells that contain electrons and designate their populations by appropriate superscripts.

19

H 1e- 1s1

# of electrons in the subshell

Hydrogen Atom Electron Configuration

principalquantum #

(shell)

Angular Momentum(shape) quantum #

(subshell)

This electronic structure accounts for the observed chemical properties of the hydrogen atom.

# of electrons in the subshell

Helium is in Group 8A, Period 1 of the Periodic Table (Atomic Number 2)

Electron Configuration For Helium

He 2e- 1s2

principalquantum #

He: the first Noble Gas; inert, unreactive; Why ? Because the n = 1 shell is filled with 2 electrons

Angular Momentum(shape) quantum #

(subshell)

Oxygen, Group 6A, Period 2; Element 8

Has 8 e-’s, the n = 1 shell is filled with 2 electrons, the other 6 electrons go into the n = 2 shell.

O 8e- 1s2 2s2 2p4

Ground State electron configuration of O:1s2 2s2 2p4

• Orbital Diagram: Shows each orbital by a box (circle) and how the electrons are placed in the orbitals with their spins specified.

• s subshells fill with electron pairs; one spin up d d h f h and one spin down in the 1s; same for the 2s

• This leaves 4 electrons for the 2p subshell.

1s2 2s2 2p4

Why did we fill the 2p subshell this way?

Many different ways to distribute the four electrons in the 3 individual p orbitals in the psubshell.

Hund's rule: For a partially filled subshell, the lowest energy arrangement will have the maximum number of unpaired electrons (orbitals

2p4

f p (with only a single electron) each with the same (parallel) spin state.

Orbital diagrams for a p subshell with 4 electronsIt doesn’t matter how these three individual porbitals are displayed, one orbital must have two electrons with paired spins and two orbitals must each contain one electron with parallel spin.

These are all equivalent orbital diagrams for a p4

subshell.

20

Orbital diagrams for a p subshell with 4 electrons

Higher energy arrangements of electrons

These are higher energy orbital diagrams for a p4 subshell.

MAGNETIC PROPERTIES OF ATOMSIf all electrons in an atom are paired in orbitals(i.e., orbitals are filled), then the magnetic moments (fields) of the electrons exactly cancel. The atom will not be affected (attracted) by a magnet. The atom is said to be diamagnetic.

e.g., Helium (He - 2e-)

1s2

diamagnetic

MAGNETIC PROPERTIES OF ATOMS

If an atom has a subshell that is partially filled, then the magnetic moments of "unpaired" electrons will add together to make the atom magnetic. The atom will be attracted by a magnet. It is said to be paramagnetic.

1s1

paramagneticH

1s2 2s2 2p4

Oxygen Paramagnetic, 2 unpaired e-’s

MAGNETIC PROPERTIES OF ATOMSPredict magnetic properties (diamagnetic or paramagnetic) from ground state electron configurations and orbital diagrams.

Consider Nitrogen with 7 electrons:

It is paramagnetic with 3 unpaired electrons.

N7e-

1s2 2s2 2p3

MAGNETIC PROPERTIES OF ATOMS

paramagnetic with 2 unpaired electrons

C6e-

1s2 2s2 2p2

p g p

Al13 e-

1s2 2s2 2p6 3s2 3p1

paramagnetic with 1 unpaired electron

MAGNETIC PROPERTIES OF ATOMS

It has no unpaired electrons, so it is diamagnetic.

Be4e-

1s2 2s2

Ne10e-

1s2 2s2 2p6

diamagnetic with no unpaired electrons

21

Figure 7.23

Ground-state electron configurations for atoms with many electrons:

Fill orbitals by energy, lowest energy orbitals (subshells) are filled according to:

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

“Aufbau principle”: German for “building up”; put electrons into the lowest energy orbitals.

A simple way to remember this order:

S t th b h ll Set up the s subshells in a column, then the psubshells in a column next to it, then the dsubshells in a column next to it, etc.

Then fill along the diagonals of the column

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

Figure 7.24

22

An Example for a Many Electron Atom

Consider Iron, Atomic Number 26 (# of protons)

26Fe 26 e-

Electron Configuration: Fill the subshells according to their energy:according to their energy:

1s2 2s2 2p6 3s2 3p6 4s2 3d6

Focus on the 3d subshell (five orbitals, partially filled subshell).

What are the magnetic properties for this electron configuration of Fe??

Fe - Paramagnetic with four unpaired electrons.

3d6

Electron configurations of the elements are listed in Table 7.3 (page 300).

Some exceptions to our guidelines-we will ignore these exceptions.

For example: Chromium (Cr-24 e-)F mp m um ( )1s2 2s2 2p6 3s2 3p6 4s2 3d4 expected 1s2 2s2 2p6 3s2 3p6 4s1 3d5 actual

For example: Copper (Cu-29 e-)1s2 2s2 2p6 3s2 3p6 4s2 3d9 expected 1s2 2s2 2p6 3s2 3p6 4s1 3d10 actual

Filled shells yield particularly stable electronconfigurations, e.g., the noble gases:

He 2e- 1s2

Ne 10e- 1s2 2s2 2p6

Ar 18e- 1s2 2s2 2p6 3s2 3p6

K 36 1 2 2 2 2 6 3 2 3 6 4 2 3d10 4 6Kr 36e- 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

We will be concerned with the outer-shellelectrons. These are called valence electrons and they control the chemical properties of the atom.

The electrons in the inner-shells are called the core electrons. The have little effect on chemical properties.

ABBREVIATED ELECTRON CONFIGURATIONS

The noble gas core will be indicated by the appropriate noble gas element symbol in bracketsfollowed by the additional electrons:

Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6 [Ar] 4s2 3d6Fe 1s 2s 2p 3s 3p 4s 3d [Ar] 4s 3d

For the halogens in Group 7A:F [He] 2s2 2p5

Cl [Ne] 3s2 3p5

Br [Ar] 4s2 3d10 4p5

Use the periodic table to determine the last subshell that is being filled:

Groups 1A and 2A s subshells

Groups 3A to 8A p subshells

Transition metals d subshells

Lanthanides f subshells

Actinides f subshells

23

Figure 7.28

Group 7A - HalogensF 2s22p5

Group 2A - Alkaline earthsBe 2s2

Mg 3s2

Why do elements in a group have similar chemical properties? Because they have similar valence shell (outer shell) electron configurations.

Cl 3s23p5

Br 4s24p5

I 5s25p5

Ca 4s2

Sr 5s2

Ba 6s2

Ra 7s2

ns2 np5 ns2