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Electron Configurations and the Periodic Table

Slide 3 / 184 The Problem with the Nuclear Atom

The nucleus of an atom is small, 1/10,000 the size of the atom. The electrons are outside the nucleus, moving freely within the vast empty atom. The nucleus is positive; the electron is negative There is an electric force, FE = kq1q2/r2, pulling the electrons towards the nucleus. There is no other force acting on the electrons;they feel a net force towards the nucleus

Why don't the electrons fall in... why doesn't the atom collapse into its nucleus?

Slide 4 / 184 The Problem with the Nuclear Model

If the Rutherford model of the atom were correct, the atom should emit energy as the orbit of the electron decays. Since the electron would speed up as it decays, the amount of energy released should be of an increasingly higher frequency. This would create what is called a continuous spectrum representing all frequencies of light.

e-

emits energy continuous spectrum

Slide 5 / 184 The Problem with the Nuclear Model

Our observations tell us the nuclear model is insufficient

• 1. Most atoms are stable and do not release energy at all

If electrons were continuously orbiting the nucleus in uniform circular motion, they would be accelerating, and accelerating charges release

• energy. This is not observed.

Slide 6 / 184 The Problem with the Nuclear Model

When energy is added to atoms, atoms do release energy in the form

• f light.

Electrons in atoms can absorb energy from collisions with photons

• r other particles and become

"excited." The excited electrons move from their initial state farther from the nucleus. Then they emit energy in the form

• f light as they return to their original

state.

Electron Absorbing Energy Electron Releasing Energy

http://imagine.gsfc.nasa.gov/docs/teachers/lessons/xray_spectra/background-atoms.html

Slide 7 / 184 The Problem with the Nuclear Model

e-

External energy added (electricity, light, etc.) light energy emitted nucleus Emission Spectrum

e-

• 2. When energized atoms do emit energy, a continuous spectrum

is not produced; instead, an emission spectrum is produced displaying emitted light at specific wavelengths and frequencies.

Slide 8 / 184

Emission Spectra and the Bohr Model

A scientist named Niels Bohr interpreted these observations and created a new model of the atom that explained the existence

• f emission spectra and provided

a framework for where the electrons can exist around the nucleus.

Slide 9 / 184

Emission Spectra and the Bohr Model

Bohr knew that the wavelengths seen in the emission spectra of hydrogen had a regular pattern. Each series was named after the scientist who observed these particular spectral lines.

Balmer Series (spectral lines in the visible and UV range) Lyman Series (spectral lines in the UV range) Paschen Series (spectral lines in the infrared range)

Slide 10 / 184

Emission Spectra and the Bohr Model

No one knew what "n" was. Bohr proposed that "n" referred to a particular orbit around the nucleus where an electron could be. Bohr proposed that electrons could orbit the nucleus, like planets

• rbit the sun...but only in

certain specific orbits. He then said that in these

• rbits, they wouldn't radiate

energy, as would be expected normally of an accelerating charge. These stable orbits would somehow violate that rule. Each orbit would correspond to a different energy level for the electron.

Slide 11 / 184

The Bohr Atom

1 2 3 4 5

n

The lowest energy level is called the ground state; the others are excited states.

Slide 12 / 184

1 An accelerating charge emits light energy. True False

Answer

Slide 13 / 184

2 When hydrogen atoms are energized by electricity, what is observed? A A continuous spectrum of light B An emission spectrum of specific colors only. C Neither a nor b

Answer

Slide 14 / 184

3 Why was the Nuclear Model insufficient? A It could not explain the existence of emission spectra B It could not account for the stability of the atom C It required the electrons to be in the nucleus and the protons in orbit around the nucleus D A and B E A, B, and C

Answer

Slide 15 / 184

4 In the Bohr model of the atom an electron in its lowest energy state A is in the ground state B is farthest from the nucleus C is in an excited state D emits energy E both a and b

Answer

Slide 16 / 184 Emission Spectra and the Bohr Model

Since atoms do not normally emit radiation, Bohr believed that the electrons existed in discrete stable orbits (n) around the nucleus which varied in energy relative to their distance from the nucleus.

n = 1 n = 2 n = 3

+ Increasing energy

Bohr was able to calculate the energy of each of these orbits. n Energy (J) 1

• 2.178 x 10-18

2

• 5.445 x 10-19

3

• 2.417 x 10-19

Slide 17 / 184 Emission Spectra and the Bohr Model

Interestingly, the energy differences between the Bohr orbits were found to correlate exactly with the energy of a particular spectral lines in the emission spectra of Hydrogen!

n = 1 n = 2 n = 3

+

Hydrogen atom

Hydrogen emission spectrum Red line wavelength ( )= 656.3 nm E = h/ E = 3.033 x 10-19 J Energy of n = 3 = -2.417 x 10-19 J Energy of n = 2 = -5.445 x 10-19 J E = (-2.417 x 10-19 J) - (-5.445 x 10-19 J) E = 3.03 x 10-19 J

EQUAL!!!

Slide 18 / 184 Emission Spectra and the Bohr Model

n = 1 n = 2 n = 3

+

Hydrogen atom

n = 4

Bohr reasoned that each spectral line was being produced by an electron "decaying" from a high energy Bohr orbit to a lower energy Bohr orbit. Since only certain frequencies of light were produced, only certain

• rbits must be possible.

Slide 19 / 184

upper lower e- upper lower e-

These possible energy states for atomic electrons were quantized –

• nly certain values were possible. The spectrum could be

explained as transitions from one level to another. Electrons would only radiate when they moved between orbits, not when they stayed in one orbit.

Emission Spectra and the Bohr Model Slide 20 / 184

According to Bohr's model, first an electron is excited from its ground state by absorbing energy.

Emission Spectra and the Bohr Model

n = 1 n = 2 n = 3

+

n = 4

photon

Slide 21 / 184 Emission Spectra and the Bohr Model

Here we see 2 separate emissions coming from the same

• electron. The electron can either go from n=3 right to n=1 or it

can go from n=3 to n=2 to n=1. Both are acceptable and both will occur. Once an electron is excited, it can take any number of routes back to its ground state, so long as it is releasing energy in discrete quantitized packets.

n = 1 n = 2 n = 3

+

n = 4 n = 1 n = 2 n = 3

+

n = 4

Slide 22 / 184

+

3 2 6 2 4

Transition

2 656 nm 486 nm 410 nm

light emitted

Review: Emission Spectrum of Hydrogen

Hydrogen atoms have one proton and one electron. The emission spectrum of hydrogen shows all of the different possible wavelengths of visible light emitted when an excited electron returns to a lower energy state.

Click here for Bohr model animation

Slide 23 / 184 Emission Spectra and the Bohr Model

The difference in energy between the orbits decreases as

• ne moves farther from the nucleus.

n = 1 n = 2 n = 3

+

Transition wavelength of spectral line produced (nm) Energy (J) 3 --> 2 656 3.03 x 10-19 2 --> 1 122 1.63 x 10-18

Slide 24 / 184 Emission Spectra and the Bohr Model

Due to the differing numbers of protons in the nucleus and number of electrons around them, each atom produces a unique emission spectrum after being energized. Since the emission spectrum

• f each element is unique, it

can be used to identify the presence of a particular element.

Slide 25 / 184 Emission Spectra and the Bohr Model

Class Question: Below are the visible wavelength emission spectra for hydrogen and iron. What difference do you notice about the two spectra and propose a reason for this difference. Iron has many more spectral lines - perhaps because it has 26 electrons to hydrogen's 1. move for answer

Emission spectrum of Hydrogen Emission spectrum of Iron

Slide 26 / 184 Flame Tests

When an excited atom emits light we see all of the spectral lines combined and only one color is visible to us. A prism or diffraction grating is needed to see the emission spectrum. light from excited atom spectral lines prism

Slide 27 / 184 Flame Tests

However, for many elements, one can identify them simply by the color produced by all of the spectral lines together. A flame test is a procedure used in chemistry to detect the presence of certain metal ions, based on each element's characteristic emission spectrum.

Slide 28 / 184

Flame Tests

Application: Fireworks make use of the fact that atoms emit visible light when excited with energy. Furthermore, they make use of the fact that each element has its own unique emission spectrum to produce the different colors you see. Calcium produces the

• range-red

color Sodium produces a yellow color

Slide 29 / 184

5 Which of the following best explains why excited atoms produce emission spectra and not continuous spectra? A Not all atoms contain enough electrons to produce a continuous spectrum B A continuous spectrum requires the movement of neutrons C Electrons can only exist in certain stable orbitals of specific energies D Electrons can exist and move anywhere around the nucleus and are not bound to a specific orbit E None of these

Answer

Slide 30 / 184

6 Which of the following is NOT true regarding the Bohr model of the atom? A Electrons could exist only in certain quantized orbits around the atom B As "n" becomes greater, the energy of the orbit is greater also C The lowest energy orbit an electron is in is referred to as the ground state D When returning from an excited state, an electron can can only move between the set Bohr orbits E All of these are true

Answer

Slide 31 / 184

7 Which of the following electron transitions would produce the highest energy spectral line? A 5 --> 4 B 3 --> 2 C 4 --> 3 D 2 --> 1 E 6 --> 5

Answer

Slide 32 / 184 Absorption vs. Emission

Since electrons can only transition between orbits of set energies atoms must absorb energy at the same frequencies at which they emit energy. As a result, monitoring which frequencies of light are absorbed can help us determine which element or molecule is present.

Slide 33 / 184

8 Does the picture below illustrate a photon

emission or absorption?

A

Emission

B

Absorbtion

C

Neither

D

Both

n = 1 n = 2 n = 3

+

n = 4

Answer

Slide 34 / 184

9 The electron in the hydrogen atom below

transitions from n=6 to n=2 and emits light with a wavelength of around 410 nm. This corresponds to which color in the visible spectrum?

A

Red

B

Orange

C

Blue

D

Violet

[This object is a pull tab]

Answer

+

656 nm 486 nm 410 nm

Hydrogen's Emission Spectrum

Slide 35 / 184

10 The electron in the hydrogen atom below

transitions from n=3 to n=2, emitting which color of light?

A

Red

B

Green

C

Blue

D

Violet

Answer

+

Slide 36 / 184

11 The emission spectrum for Chlorine is shown

• below. Which of the following represents

Chlorine's corresponding absorption spectrum?

A B C Answer

Slide 37 / 184

12 One atom of hydrogen can produce the entire hydrogen emission spectrum. True False

Answer

Slide 38 / 184 The Bohr Atom and Atomic Radii

Another piece of evidence to support the Bohr model was that it was also able to accurately predict the size of an atom using the Coulomb force and his orbit concept. First, an electron is held in orbit by the Coulomb force:

Slide 39 / 184 The Bohr Atom

Using the Coulomb force, we can calculate the radii

• f the orbits. These matched the sizes of known

atoms very well. Where: rn = radius of Bohr Orbit "n" Z = nuclear charge r1 = radius of Bohr Orbit 1

Slide 40 / 184 The Bohr Atom

The radii of the orbits of a hydrogen atom are given by the below formula, with the smallest orbit, r1 = 0.53 x 10-10 m. n = 1, 2, 3, 4, ....

n2r1

Z

rn =

Notice that the orbits grow in size as the square of n, so they get much larger as n increases. (for hydrogen, Z = 1)

Slide 41 / 184

13 The radius of the orbit for the third excited

state (n=4) of hydrogen is ______r1.

Answer

Slide 42 / 184

14 The radius of the orbit for the fifth excited

state (n=6) of hydrogen is ____ x 10-10 m. r1 = 0.50 x 10-10 m

Answer

Slide 43 / 184 de Broglie’s Hypothesis and the Bohr Model

de Broglie's wave theory of matter also explained Bohr's orbits

• well. As long as the wavelength of an electron in orbit was the

same as the circumference of the orbit, it would not radiate and move into the nucleus.

de Broglie proposed electrons in specific orbits produced standing waves at specific wavelengths, frequencies, and energies.

Slide 44 / 184

Quantum Physics

While a big step forward, Bohr's model was only useful in predicting the frequency of spectral line for atoms that had one electron, like hydrogen or certain ionized atoms. The idea that the electron was a particle in orbit around the nucleus, but with wavelike properties that only allowed certain orbits, worked only for hydrogen. Semi-classical explanations failed except for hydrogen. It turned out that only a lucky chance let it work even in that case.

Slide 45 / 184 Quantum Mechanics

Our goal was to explain why electrons in an atom don't fall into the nucleus. An electron, as a charged particle, would fall in because of Newton's Second Law. ΣF = ma But electrons, in atoms, aren't particles, they're waves. Waves don't follow Newton's Second Law. Schrodinger had to invent a new equation for wave mechanics. Hψ = Eψ

slide to reveal new equation

Slide 46 / 184 Electrons: wave or particle or both?

Recall that electrons can behave as both a particle and a wave. With this in mind, to find where they are in the atom we have to consider both the particle and wave nature of the electron. Bohr's simple atomic model was not sufficient to explain the position of electrons, so a new model was needed.

Slide 47 / 184 Quantum Mechanics – A New Theory

Quantum mechanics is a branch of physics which provides a mathematical description of wave-particle duality, and successfully explains the following 2 ideas: (1) the energy states in complex atoms and molecules (2) the relative brightness of spectral lines

Slide 48 / 184 Quantum Mechanics – A New Theory

It is widely accepted as being the fundamental theory underlying all physical processes. Quantum mechanics is essential to understanding atoms and molecules, but can also have effects on larger scales.

Slide 49 / 184 The Wave Function and Its Interpretation

An electromagnetic (light) wave is made of oscillating electric and magnetic fields. What is oscillating in an electron or matter wave? The wave function, Ψ (psi) describes the state and behavior of an electron. The two fields of the wave are noted in blue and red in this animation. Each wave frequency is proportional to the possible energy level of the

• scillator.

Slide 50 / 184 Interpretation of the Wave Function (Ψ)

The square of the wave function at any point is proportional to the number of electrons expected to be found there.

∞

Ψ2 # electrons For a single electron, the wave function is the probability of finding the electron at that point. Ψ = Probability of finding electron

Slide 51 / 184 The Double-Slit Experiment

Recall the interference pattern

• bserved after many electrons

have gone through the slits. If we send the electrons through one at a time, we cannot predict the path any single electron will take, but we can predict the

• verall distribution.

Light or Electrons Intensity

• n screen

Slide 52 / 184

This series of photos was made by electrons being fired one at a time through two slits. Each exposure was made after a slightly longer time. The same pattern emerges as was found by

• light. The only explanation is that each

individual electron behaves like a wave as it passes through both slits. But each electron must be a particle when it strikes the film, or its wouldn't make one dot

• n the film, it would be spread out.

Young's Double Slit Experiment

This one picture showed that matter is both a wave and a particle.

Slide 53 / 184

15 The probability of finding an electron at a specific location is directly proportional to:

A

its energy. B its momentum.

C

its wave function. D the square of its wave function.

Answer

Slide 54 / 184 The Heisenberg Uncertainty Principle

Quantum mechanics tells us there are inherent limits to measurement. This is not because of the limits of our instruments, rather it is due to the wave-particle duality, and to the interaction between the observing equipment and the object being

• bserved.

With this in mind, in 1926 a man named Werner Heisenberg proposed what's known as the Heisenberg Uncertainty Principle.

Slide 55 / 184

The Heisenberg Uncertainty Principle

Imagine trying to see an electron with a powerful microscope. At least one photon must scatter off the electron and enter the microscope. However, in doing so, it will transfer some of its momentum to the electron.

Slide 56 / 184

Alternate explanation of the Uncertainty Principle

Think of electrons like small ball bearings rolling around in a large warehouse. We cannot directly see (or hear) electrons because they are so small, so we will imagine the warehouse is fully dark and we are wearing ear plugs. In this experiment, in order to find an electron you are given a stick to skim around on the floor to "feel" for the electrons (since we can "feel" or observe an electrons effect).

Slide 57 / 184

Ball Bearing Experiment

What happens to the ball bearings' position once you locate it by tapping it with the stick? If we ignore friction, and even allow our hypothetical electron to fly around the room in 3 dimensions (like electrons actually do) could we ever really know where the ball bearing is EXACTLY? Of course not! Its the same thing with electrons. They are so small that the very act of

• bserving their position changes their position.

It is a vicious cycle!

Slide 58 / 184 The Heisenberg Uncertainty Principle

The uncertainty in the momentum of the electron is taken to be the momentum of the photon; meaning it could transfer none of its momentum or all of is momentum

in our example, the stick could hit the ball bearing and bounce off, or cause the ball bearing to bounce off

In addition, the position can only be measured to within about

• ne wavelength of the photon

in our example, we might miss the ball bearing completely

Slide 59 / 184 The Heisenberg Uncertainty Principle

Combining, we find the combination of uncertainties:

( x) ( px ) h

This specifically is the Heisenberg uncertainty principle. It tells us that the position and momentum of a particle cannot simultaneously be measured with precision.

Slide 60 / 184

The Heisenberg Uncertainty Principle

This can also be written as the relationship between the uncertainty in time and the uncertainty in energy: This says that if an energy state only lasts for a limited time, its energy will be uncertain. It also says that conservation of energy can be violated if the time is short enough.

( E) ( t) h

Slide 61 / 184

16 The reason the position of a particle cannot be specified with infinite precision is the:

A

exclusion principle.

B

uncertainty principle.

C

photoelectric effect.

D

principle of relativity. Answer

Slide 62 / 184

17 If the accuracy in measuring the position of a particle increases, the accuracy in measuring its velocity will:

A

increase.

B

decrease.

C

remain the same.

D

be uncertain. Answer

Slide 63 / 184

18 If the accuracy in measuring the velocity of a particle increases, the accuracy in measuring its position will:

A

increase.

B

decrease.

C

remain the same.

D

be uncertain. Answer

Slide 64 / 184 Philosophic Implications: Probability versus Determinism

As you know, the world of Newtonian mechanics is a deterministic one. If you know the forces on an object and its initial velocity, you can predict where it will go. Quantum mechanics is very different. You can predict what most electrons will do on average, but you can have no idea what any individual electron will do.

Slide 65 / 184 Classical versus Quantum Mechanics

In classical physics, predictions about how objects respond to forces are based on Newton's Second Law: ΣF = ma In quantum physics, this no longer works; predictions are based on Schrödinger's Wave Equation. Hψ = Eψ

Slide 66 / 184

Schrödinger's Wave Equation

Solving this equation is well beyond this course. And only probabilities of outcomes can be determined…you cannot specifically determine what will happen in each case. However, this equation has been solved for many specific cases and we will be using those solutions to understand atoms, molecules, and chemical bonds.

Hψ = Eψ

Slide 67 / 184

Schrödinger and his cat?

Erwin Schrödinger received the Nobel Prize in Physics in 1933 for the development of the Schrödinger Equation. Additionally he is known for his famous thought experiment where he applied quantum mechanics to everyday objects... specifically a cat.

click here for a short explanation

• f "Schrodinger's Cat"

Slide 68 / 184

19 Quantum mechanics provides a mathematical definition for the:

A

wave-like properties of electrons only. B particle-like properties of electrons only

C

classic Newtonian forces that govern atoms D the wave-particle duality of electrons

Answer

Slide 69 / 184 Quantum-Mechanical View of Atoms

Since we cannot say exactly where an electron is, the Bohr picture of the atom, with its electrons in neat orbits, cannot be correct. Quantum theory describes an electron probability distribution; this figure shows the distribution for the ground state of hydrogen. In this picture, the probability of finding an electron somewhere is represented by the density of dots at that location.

Slide 70 / 184

Quantum Numbers

Solutions to Schrodinger's Wave Equation take the form of sets of

• numbers. There are four different quantum numbers: n, l, ml, ms

needed to specify the state or probable location of an electron in an atom.

n = 1 n = 2 n = 3

+

n = 4

n = principal l = angular

X Y Z

ml = magnetic ms = spin energy level/ distance from nucleus shape

• f orbital
• rientation of
• rbital in space

direction of electron spin

+

• Slide 71 / 184

(n): Principal Quantum Number

An orbital is a region of space where an electron is most likely to be found. The principal quantum number, n, describes the energy level of the orbital. The values of n are integers greater than or equal to 1: n ≥ 1 In general, the larger the value of n, the farther from the nucleus the electron should be found

n = 1

n = 2 n = 3

+

n = 4

Slide 72 / 184

20 The principal quantum number, n, determines the ____________ of the orbital.

A

Orientation

B

Energy

C

Shape

D

Capacity Answer

Slide 73 / 184

21 As n increases, the orbital energy _________ .

A

Increases

B

Decreases

C

Remains constant

D

Increases then decreases Answer

Slide 74 / 184

Quantum number l designates the shape of the orbital. There are four shapes of orbitals: s,p,d,f Quantum number ml designates the orientation of the

• rbital in space.

(l): Angular Quantum Number (ml): Magnetic Quantum Number

Each orbital region has a very specific shape based on the energy of the electrons occupying them and a specific

• rientation in space.

Slide 75 / 184 Electron Orbital Shape and Orientations

This quantum number defines the shape of the orbital, and gives the angular momentum.

http://chemwiki.ucdavis.edu/@api/deki/files/4826/=Single_electron_orbitals.jpg

Slide 76 / 184 s Orbitals

s orbitals are spherical in shape.The radius of the sphere increases with the value of n. If you are looking for an electron in an s orbital, the direction you look in doesn't really matter, they have only one orientation in space.

1 2

3

1 2

3

If l = s shape ml = 1 orientation 1 orbital per energy level

Slide 77 / 184

p Orbitals

p orbitals have two lobes with a node between them. For p orbitals, the amount of electron density and the probability of finding an electron depends on both the distance from the center of the atom, as well as the direction. The p orbital has 3 possible arrangements in space, so it can have 3 possible orbitals.

High probability of finding an electron Low probability of finding an electron

l = p shape ml = 3 orientations 3 orbitals per energy level

Slide 78 / 184 d orbitals

d orbitals have more complex shapes. There are 5 possible orientations in space, so there are 5 possible d orbitals.

l = d shape ml = 5 orientations 5 orbitals per energy level

Slide 79 / 184 f Orbitals

There are 7 possible f orbitals.

l = f shape ml = 7 orientations 7 orbitals per energy level

Slide 80 / 184

22 The quantum number, l, determines the ____________ of the orbital.

A

Orientation

B

Energy

C

Shape

D

Capacity Answer

Slide 81 / 184

23 The magnetic quantum number, ml,

determines the ____________ of the orbital.

A

Orientation

B

Energy

C

Shape

D

Capacity Answer

Slide 82 / 184

24 A(n) ___ orbital is lobe-shaped A s

B

p

C

d

D

f Answer

Slide 83 / 184

25 An s orbital has ______ possible

• rientations in space.

A

1

B

3

C

5

D

7 Answer

Slide 84 / 184

26 An f orbital has ______ possible

• rientations in space.

A

1

B

3

C

5

D

7 Answer

Slide 85 / 184 Spin Quantum Number, ms

In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy. This led to a fourth quantum number, the spin quantum number, m

s.

Slide 86 / 184 Spin Quantum Number, ms

The “spin” of an electron describes its magnetic field, which affects its energy. The spin quantum number can be positive or negative. This implies that electrons are in some way able to pair up, even though they repel each other due to the electromagnetic force. Each orbital can therefore hold a maximum of 2 electrons.

+ spin

• spin

Slide 87 / 184

27 The spin quantum number, m

s

A

can only have two values

B relates to the spin of the electron

C

relates to the spin of the atom

D

Both A & B

E

A, B & C Answer

Slide 88 / 184

28 Electrons within the same

• rbital must have

A

the same spin

B no spin

C

• pposite spins

D

electrons cannot occupy the same orbital

Answer

Slide 89 / 184

The Four Quantum Numbers

The quantum state of an electron is specified by the four quantum numbers; no two electrons can have the same set of quantum numbers.

n, Principal quantum number designates energy or shell level l, angular quantum number designates orbital shape: s,p,d,f ml, magnetic quantum number designates orbital orientation ms, designates electron spin +

• Click here for a Review Video

Slide 90 / 184

Remember: an orbital is a location with a high probability of finding an electron. Shell - refers to the first quantum number, n Subshell - refers to the second quantum number, l and the specific type of orbitals (s,p,d, or f) within a given shell Orbital - refers to the third quantum number, ml

Quantum Number Vocabulary

Slide 91 / 184

Energy Levels and Subshells

Some combinations of Quantum Numbers are impossible: If n = 1, an electron can only occupy an s subshell. If n = 2, an electron can only occupy s or p subshells. If n = 3, an electron can only occupy s, p, or d subshells If n = 4 an electron can occupy s, p, d, or f subshells

Slide 92 / 184

Quantum Numbers Subshells

Orbitals with the same value of n form a shell. Different orbital types within a shell are subshells.

n subshell # of orbitals total # total #

• f orbitals of electrons

1 1s 1 1 2 2 2s 1 2p 3 4 8 3 3s 1 3p 3 3d 5 9 18 4 4s 1 4p 3 4d 5 4f 7 16 32

Slide 93 / 184

29 If n = 1 an electron can occupy which of the subshells?

A 1s B

2s

C 2p D 3s

Answer

Slide 94 / 184

30 An s orbital can hold a maximum of ___ electrons

A

1

B 2

C

4

D

infinity Answer

Slide 95 / 184

31 What is the maximum number of electrons that can occupy the f orbital?

A

10

B 14

C

18

D 22

Answer

Slide 96 / 184

32 How many possible sets of quantum numbers

• r electron states are there in the 4d orbital?

A

2

B

8 C

10

D 14

Answer

Slide 97 / 184

33 An electron is in the 6f state. Determine the principal quantum number.

A

3

B

5

C

6

D

14

Answer

Slide 98 / 184

34 An electron is in the 6d state. How many electrons are allowed in this state?

A

6

B

7

C

10

D

14

Answer

Slide 99 / 184

35 How many electrons will fit into a 4f orbital?

A

3

B

7 C 14

D 4

Answer

Slide 100 / 184

36 An electron is in the 6f state. How many electrons are allowed in this state?

A

6

B

7

C

10

D

14

Answer

Slide 101 / 184

We will be learning how to list all of the electrons in an atom based on the orbital each electron falls in. To do this we will also need to know all about the order of the orbitals based on the energy required to put electrons into them. This will affect certain elements and will cause some other elements to be exceptions to the general order of orbitals.

Electron Inventory Slide 102 / 184

In order to understand each topic we will need to understand the topic that comes before it in the triangle diagram below. This might be a bit confusing, but keep in mind, the entire concept will not come together until all the pieces are there.

Electron Inventory

Order of Orbitals Energy of Orbitals Exceptions to the Order

Slide 103 / 184 Energies of Orbitals

For a one-electron hydrogen atom, orbitals on the same energy level have the same energy. Degenerate is the term referring to the fact that the energies are the same.

Slide 104 / 184 Energies of Orbitals

As the number of electrons increases, though, so does the repulsion between them. Complex atoms contain more than one electron, so the interaction between electrons must be accounted for in the energy levels. This means that the energy depends

• n both n and l.

Therefore, in many-electron atoms,

• rbitals on the same energy level are

no longer degenerate.

Slide 105 / 184 Energies of Orbitals

All these graphs show how energy depends on n and l, just in different depictions

Slide 106 / 184

1 2 3 4 5 6 7

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

Energy

Energies of Orbitals

Take note of how close in energy the d and f

• rbitals are to the s
• rbitals of later energy

levels. This will have an impact on exceptions to several rules we will discuss a little later this chapter.

Slide 107 / 184 Orbital Diagrams

Each box in the diagram represents one orbital. Orbitals on the same sublevel are drawn together. Arrows represent the electrons. The direction of the arrow represents the relative spin (+ or - ) of the electron.

8O

1s 2s 2p Slide 108 / 184 3 Rules for Filling Electron Orbitals

Aufbau Principle Electrons are added one at a time to the lowest energy

• rbitals available until all the electrons of the atom have

been accounted for. Pauli Exclusion Principle An orbital can hold a maximum of two electrons. To

• ccupy the same orbital, two electrons must spin in
• pposite directions.

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

Slide 109 / 184

Aufbau takes its name from a German word meaning "building up". Developed in the 1920s by Bohr and Pauli and states that Electrons fill the lowest energy

• rbitals first.

Aufbau Principle

Imagine it in terms of Lazy Tenants - tenants in a multistory building fill in from the ground level up, so they don't have to walk up stairs

Slide 110 / 184

No two electrons in the same atom can have exactly the same energy.

Pauli Exclusion Principle

The quantum state is specified by the four quantum numbers; no two electrons can have the same set of quantum numbers (ms = + or -)

1s2 2s2 2p1

correct

1s2 2s2 2p1

incorrect

Slide 111 / 184

Think about the Empty Bus Seat Rule. People will not sit next to each other on a bus until all the seats are taken up

Hund’s Rule

Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

1s2 2s2 2p2

correct

1s2 2s2 2p2

incorrect

Slide 112 / 184

Energy Level Diagram

Fill in the Energy Level Diagram for Magnesium, Mg:

Slide 113 / 184

Energy Level Diagram

Fill in the Energy Level Diagram for Chlorine, Cl:

Slide 114 / 184

Energy Level Diagram

Fill in the Energy Level Diagram for Iron, Fe:

Slide 115 / 184 Electron Configurations Slide 116 / 184 Electron Configurations

Electron configurations show the distribution of all electrons in an atom. Each component consists of: A number denoting the energy level,

Slide 117 / 184 Electron Configurations

Electron configurations show the distribution of all electrons in an atom. Each component consists of: A number denoting the energy level, A letter denoting the type of orbital,

Slide 118 / 184 Electron Configurations

Electron configurations show the distribution of all electrons in an atom. Each component consists of: A number denoting the energy level, A letter denoting the type of orbital, and A superscript denoting the number of electrons in those orbitals.

Slide 119 / 184

Aufbau Mneumonic

Slide 120 / 184 Electron Configuration of Sodium

For example, here is the ground-state configuration of sodium:

1s2 2s2 2p6 3s1

Na

23 11

Sodium Atom (All of the superscript numbers add up to the total number of electrons)

Slide 121 / 184 Energy Level Diagram

Electron Configuration Write the Ground State Electron Configuration for Magnesium, Mg:

Slide 122 / 184 Energy Level Diagram

Electron Configuration Write the Ground State Electron Configuration for Chlorine, Cl:

Slide 123 / 184 Energy Level Diagram

Electron Configuration Write the Ground State Electron Configuration for Iron, Fe:

Slide 124 / 184 Energy Level Diagram - Excited State

In a sodium-vapor lamp electrons in sodium atoms are excited to the 3p level by an electrical discharge and emit yellow light as they return to the ground state. Na Excited State Energy Level Diagram

Slide 125 / 184

Energy of Orbitals

Click here to view an Interactive Periodic Table that shows orbitals for each Element Click here for an electron orbital game.

Slide 126 / 184

37 A neutral atom has the following electron configuration: 1s22s22p63s23p2. What element is this?

A carbon B nitrogen C silicon D germanium

Answer

Slide 127 / 184

38 A neutral atom has an electron configuration of 1s22s22p63s23p2. What is its atomic number?

A

5

B 11

C 14

D 20

Answer

Slide 128 / 184

39 A neutral atom has an electron configuration

• f 1s22s22p6. If a neutral atom gains one

additional electron, what is the ground state configuration?

A

1s22s22p63s1

B

1s22s22p7

C

1s22s32p6

D

none of the given answers Answer

Slide 129 / 184

40 Which of the following represents an excited state electron configuration for Sodium (Na)?

A

1s22s22p63s1

B

1s22s22p7

C

1s22s22p63p1

D

none of the given answers Answer

Slide 130 / 184 Shorthand Configurations

Elements on the far right column of the periodic table have their subshells entirely filled. This group of elements is referred to as the "Noble Gases." Noble Gas elements are used to write shortened electron configurations. Noble Gases

Slide 131 / 184 Shorthand Configurations

To write a Shorthand Configuration for an element: (1) Write the Symbol of the Noble Gas element from the row before it in brackets [ ]. (2) Add the remaining electrons by starting at the s orbital of the row that the element is in until the configuration is complete.

Slide 132 / 184 Shorthand Configurations

Electron Configuration: 1s22s22p63s1 Shorthand Configuration: [Ne] 3s1

Neon's electron configuration

Example: Sodium (Na)

Slide 133 / 184 Fill in Shorthand Configurations

Slide for Answers

Element Shorthand Configuration

Slide 134 / 184

41 What is the electron configuration for Li ?

A

1s3 B 1s1 2s2

C

1s2 2s1

D

1s2 1p1

Answer

Slide 135 / 184

42 Which of the following would be the correct electron configuration for Mg?

A 1s22s23s23p64s2 B 1s22s23s23p6 C 1s22s22p6 D 1s22s22p63s2 E None of these Answer

Slide 136 / 184

43 Which of the following would be the correct electron configuration for a Mg2+ ion?

A 1s22s23s23p64s2 B 1s22s23s23p6 C 1s22s22p6 D 1s22s22p63s2 E None of these Answer

Slide 137 / 184

44 Which of the following would be the correct electron configuration for a Cl- ion?

A 1s22s23s23p6 B 1s22s23s23p5 C 1s22s22p6 D 1s22s22p63s1 E None of these Answer

Slide 138 / 184

45 What would be the expected "shorthand" electron configuration for Sulfur (S)?

A [He]3s23p4 B [Ar]3s24p4 C [Ne]3s23p3 D [Ne]3s23p4 E [Ne]3d1 Answer

Slide 139 / 184

46 What would be the expected "shorthand" electron configuration for vanadium (V) ?

A [He]4s23d1 B [Ar]4s23d104p1 C [Ar]4s23d3 D [Kr]4s23d1 E [Ca]3d1 Answer

Slide 140 / 184 Stability

When the elements were studied scientists noticed that some of them do not react in certain situations in which others do. These elements were labeled "stable" because they did not change easily. When these stable elements were grouped together, it was noted that periodically, there were patterns in the occurrence of stable elements. Today we recognize that this difference in stability is due to electron configurations.

Slide 141 / 184

There are two methods for labeling the groups, the older method shown in black on the top and the newer method shown in blue on the bottom.

1A 2A 8A 1 2 18 3A 4A 5A 6A 7A 13 14 15 16 17 8B 3B 4B 5B 6B 7B 1B 2B 3 4 5 6 7 8 9 10 11 12

}

Group Numbers

Slide 142 / 184 Stability

Elements of varying stability fall into one of 3 categories. The most stable atoms have completely full energy levels. ~Full Energy Level ~Full Sublevel (s, p, d, f) ~Half Full Sublevel ( d

5, f7)

1 2 3 4 5 6 7 6 7

Slide 143 / 184 Stability

Next in order of stability are elements with full sublevels. ~Full Energy Level ~Full Sublevel (s, p, d, f) ~Half Full Sublevel ( d

5, f7)

1 2 3 4 5 6 7 6 7

Slide 144 / 184 Stability

Finally, the elements with half full sublevels are also stable, but not as stable as elements with fully energy levels or sublevels. ~Full Energy Level ~Full Sublevel (s, p, d, f) ~Half Full Sublevel ( d 5, f7) 1 2 3 4 5 6 7 6 7

Slide 145 / 184 Electron Configuration Exceptions

You should know the basic exceptions in the d- and f-sublevels. These fall in the circled areas on the table below. 1 2 3 4 5 6 7 6 7

Slide 146 / 184

Chromium Expect: [Ar] 4s2 3d4 Actually: [Ar] 4s1 3d5 For some elements, in order to at least get a half full d sublevel, electrons from an s sublevel will move to a d sublevel. To see why this can happen we need to examine how "close" d and s sublevels are.

Electron Configuration Exceptions

1 2 3 4 5 6 7 6 7

Cr

Slide 147 / 184

1 2 3 4 5 6 7

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

4f 5d 6p 5f

7s 6d 7p

6f 7d 7f

Energy

Energies of Orbitals

Because of how close the f and d orbitals are to the s

• rbitals an electron can easily

move from the s orbital (leaving it half full) to the f or d orbital, causing them to also be half full. (It's kind of like borrowing a cup of sugar from a neighbor. You'd only borrow it from someone you were close to and only if you needed it.)

Slide 148 / 184

Copper Expected: [Ar] 4s2 3d9 Actual: [Ar] 4s1 3d10 Copper gains stability when an electron from the 4s

• rbital fills the 3d orbital.

Electron Configuration Exceptions

1 2 3 4 5 6 7 6 7

Cu

Slide 149 / 184

47 The highlighted elements below are in the

A s block B d block C p block D f block Answer

Slide 150 / 184

48 The highlighted elements below are in the

A s block B d block C p block D f block Answer

Slide 151 / 184

49 The highlighted elements below are in the

A s block B d block C p block D f block Answer

Slide 152 / 184

50 The electron configuration for Copper (Cu) is

A [Ar] 4s24d9 B [Ar] 4s14d9 C [Cr] 4s23d9 D [Ar] 4s23d9 Answer E [Ar] 4s13d10

Slide 153 / 184

51 What would be the expected "shorthand" electron configuration for Silver (Ag)?

A [Kr]5s25d9 B [Ar]5s24d9 C [Ar]5s14d10 D [Kr]5s24d9 E [Kr]5s14d10 Answer

Slide 154 / 184

52 What would be the expected "shorthand" electron configuration for Molybdenum (Mb)?

A [Kr]5s25d4 B [Ar]5s24d4 C [Ar]5s14d5 D [Kr]5s14d5 E [Kr]5s24d4 Answer

Slide 155 / 184 The Periodic Table

Now that we know where (or approximately where) to find the parts of atoms, we can start to understand how these factors all come together to affect how we view the elements. We can look at them as individual yet interacting chemicals, and we are able to group them based, not only on the properties they present when in isolation, but also the properties they reveal when exposed to other elements or compounds.

Slide 156 / 184 History of the Periodic Table

Dmitri Mendeleev, building on the ideas from chemists before him, developed the modern periodic table. He argued that element properties are periodic functions of their atomic weights. We now know that element properties are periodic functions of their atomic number. By elemental properties, we are describing both physical and chemical properties. Atoms are listed on the periodic table in rows, based on number of protons.

Slide 157 / 184

A periodic table usually has the following information:

Information on the Periodic Table

NOTE: A periodic table may have more information or less information, depending on the publisher and intended use. Atomic Number - the number of protons in that particular atom Atomic Mass - the average atomic mass for that atom Name of Atom Element Symbol - the one or two letters designating the atom

Slide 158 / 184 Periodic Table

The periodic table is made of rows and columns: Rows in the periodic table are called Periods. Columns in the periodic table are called Groups. Groups are sometimes referred to as Families, but "groups" is more traditional.

Slide 159 / 184

periods groups

1 2 3 4 5 6 7

* ** ** *

6 7

Slide 160 / 184 Periodic Table

The periodic table is "periodic" because of certain trends that are seen in the elements. Properties of elements are functions of their atomic number. Elements from the same group have similar physical and chemical properties. Atoms are listed on the periodic table in rows, based on number

• f protons, which is equal to the number of electrons in a neutral

atom.

Slide 161 / 184

53 What is the atomic number for the element in period 3, group 16?

Answer

Slide 162 / 184

54 What is the atomic number for the element in period 5, group 3?

Answer

Slide 163 / 184 Special Groups

Some groups have distinctive properties and are given special names.

Alkali Metals Alkaline Earth Metals Halogens Noble Gases Transition Metals

Slide 164 / 184 Groups of Elements

Enjoy Tom Lehrer's Famous Element Song!

Slide 165 / 184

Alkali Metals

Group 1 Alkali Metals (very reactive metals)

Slide 166 / 184

Alkaline Earth Metals

Group 2 Alkaline Earth Metals (reactive metals)

Slide 167 / 184

Transition Metals

Groups 3 - 12 Transition Metals (low reactivity, typical metals)

Slide 168 / 184

Group 16 Oxygen Family (elements of fire)

Slide 169 / 184

Halogens

Group 17 Halogens (highly reactive, nonmetals)

Slide 170 / 184

Group 18 Noble Gases (nearly inert)

Noble Gases

Slide 171 / 184

Alkali Metals Alkaline Earth Metals Halogens Noble Gases Transition Metals

Slide 172 / 184 Metals, Nonmetals, and Metalloids

The periodic table can be also divided into metals (blue) and nonmetals (yellow) . A few elements retain some of the properties of metals and nonmetals, they are called metalloids (pink).

As B Si Te Ge Sb ?

metals nonmetals metalloids Slide 173 / 184 Diatomic Elements

Seven elements in the periodic table are always diatomic. In elemental form, they are always seen as two atoms bonded together. H2 , O2, N2 , Cl2 , Br2 , I2 , F2

H

O N Cl Br I F

Slide 174 / 184

Since the families are based on reactivates, and next, how something reacts is based off of how its electrons are

• arranged. . .

. . . we now know that elements in the same family have very similar electron configurations

Electron Configuration

Alkali Metals Alkaline Earth Metals Halogens Noble Gases Transition Metals

Slide 175 / 184 Group names

Noble Gases - Group 18, s2p6 ending Have a full outermost shell Halogens - Group 17, s2p5 ending Highly reactive, need one electron to have a full outer shell. Alkali Metals - Group 1, s1 ending Very reactive Alkaline Earth Metals - Group 2, s2 ending Reactive Transition Metals (d-block) - Groups 3 - 12 somewhat reactive, typical metals, ns2, (n-1)d ending Inner transition metals ( f -block) - the bottom two rows somewhat reactive and radioactive , ns2, (n-2)f ending

Slide 176 / 184

55 The elements in the periodic table that have completely filled shells or subshells are referred to as:

A

noble gases.

B

halogens.

C

alkali metals.

D

transition elements.

Answer

Slide 177 / 184

56 The elements in the periodic table which lack

• ne electron from a filled shell are referred to

as: A

noble gases.

B

halogens.

C

alkali metals.

D

transition elements.

Answer

Slide 178 / 184

57 The elements in the periodic table which have a single outer s electron are referred to as:

A

noble gases.

B

halogens.

C

alkali metals.

D

transition elements.

Answer

Slide 179 / 184

58 Which of the following represents an electron configuration of an alkaline earth metal?

A [He]2s1 B [Ne]3s23p6 C [Ar]4s23d2 D [Kr]5s24d105p4 E [Xe]6s2 Answer

Slide 180 / 184

59 Which of the following represents an electron configuration of a Halogen?

A [He]2s1 B [Ne]3s23p5 C [Ar]4s23d2 D [Kr]5s24d105p4 E [Xe]6s2 Answer

Slide 181 / 184

60 The electron configuration [Ar]4s23d5 belongs in which group of the periodic table

A Alkali Metals B Alkaline Earth Metals C Transition Metals D Halogens E Noble Gases Answer

Slide 182 / 184

61 The electron configuration ending ns2p6 belongs in which group of the periodic table

A Alkali Metals B Alkaline Earth Metals C Transition Metals D Halogens E Noble Gases Answer

Slide 183 / 184 Looking back at the Periodic Table of the Elements

Atoms with the same number of electrons in their outer shells or same outer electron configuration, have similar chemical behavior. They appear in the same column of the periodic table. The periodic table of elements can be grouped into blocks based

• n electron configuration of the atoms. s, p, d, and f blocks will

have the last electron in the atom filling into these subshells respectively. The elements with full or half full outer subshells are the most stable .

Slide 184 / 184