chapter 5 periodicity and atomic structure

8/3/2019 Chapter 5 Periodicity and Atomic Structure

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PerodicityPerodicityand Atomic Structureand Atomic Structure

Chapter 5

Chapter 5 2

Mendeleevs Periodic Table In the 1869,Dmitri Mendeleevproposed that

the properties of the chemical elements repeatat regular intervals when arranged in order ofincreasingatomic mass.

Mendeleev is the architect of the modernperiodic table.


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Chapter 5 3

Prediction of New Elements

Mendeleev noticed that there appeared to be some elementsmissing from the periodic table so he left spots for undiscovered

elements.

Based on his periodic table, he accurately predicted the properties

of some of these the unknown elements before their discovery.

Chapter 5 4

Periodic Trends The arrangement of the periodic table means that

the physical properties of the elements follow a

regular pattern.

Some trends include:

Atomic Radius (end of Chapter 5)

Ionization Energy (Chapter 6) Electron Affinity (Chapter 6)

Electronegativity (Chapter 7)


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Chapter 5 5

The Structure of Atoms

Plum Pudding Model Tootsie Pop Model

Chapter 5 6

The Dual Nature of Light:The Particle and The Wave

A great animation of the idea of light as waves and particles can be found here:

http://video.google.com/videoplay?docid=-4237751840526284618&q=quantum

From the time of the ancient

Greeks, people have thought of

light as a stream of tiny

particles - like marbles or

billiard balls

Thomas Young (in 1807)

performed the now classic

double slit experiment to testthis theory.


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Chapter 5 7

As evidenced by the double slit

experiment, light travels through space as

a wave, similar to an ocean wave.

Wavelength () is the distance light

travels in one cycle.

Frequency () is the number of wave

cycles completed each second.

Amplitude is the height measured from

the center of the wave. The square of this

value gives theIntensity.

Velocity (c) =

The Dual Nature of Light: The Wave

Light has a constant velocity (c) of3.00 108 m/s.

Chapter 5 8

Frequency - Wavelength

The red light in a laser pointer comes from a

diode laser that has a wavelength of about 632

nm. What is the frequency of the light?


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Chapter 5 9

The Dual Nature of Light: The Particle

In 1900, Max Planck proposedthat radiant energy is notcontinuous, but is emitted assmall bundles of energy

This is the quantum concept.

Planck determined that the energy of an emitted bundle is directlyproportional to the frequency of the emitted light times a constant (h)

E= h =

hc h = 6.626 10 34Js

Chapter 5 10

The Photoelectric Effect The Photoelectric Effect shows that when light is

shined on the surface of a metal, electrons are ejected

and a current can be detected.

Albert Einstein used Plancks idea ofquanta to

explain this phenomona.

If the light were waves only, then the energy of that

radiation would depend only on the intensity of the

light.

Einstein hypothesized that electrons are only ejected

if the frequency of the light exceeds a threshold

value specific to the metal, regardless of light

intensity

Even at low light intensities, electrons are ejected

immediately if the frequency exceeds the threshold

Millikan tested Einsteins hypothesis in 1914 and

proved that is was correct

Einstein wins the Nobel Prize! Yea!

Figure above from http://en.wikipedia.org/wiki/Photoelectric_Effect


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Chapter 5 11

Energy of Emission

For red light with a wavelength of about 632

nm, what is the energy of a single photon and

one mole of photons?

Chapter 5 12

Louis de Broglie suggested that if light can behave

like matter (particles) then matter can behave as light.

This concept is called wave particle duality.

For Light

=

h

mc

For a Particle

=

h

mv

Wave Particle Duality

E = mc2 E = h

h = 6.626 x 10-34 kg m2 / s h = 6.626 x 10-34 J s


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Chapter 5 13

Wave Particle Duality

How fast must an electron be moving if it has a

de Broglie wavelength of 551 nm?

How fast must an electron be moving if it has a

de Broglie wavelength of 631 nm?

me = 9.109 x 1031 kg

Chapter 5 14

The Electromagnetic Spectrum The complete electromagnetic spectrum (all possible

wavelengths and frequencies) is an un-interruptedband, orcontinuous spectrum.

The radiant energy spectrum includes most types ofradiation, most of which are invisible to the humaneye.

ROY G BIV


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Chapter 5 15

The Atomic Spectra

When a particular element isheated and the light emitted isfocused and passed through aprism, the resulting breakdownresults in anon-continuousspectrum or adiscrete linespectrum

This spectrum is called theelementsatomic spectrum

The discrete lines indicate alight is emitted or absorbedin a series of discretefrequencies rather thancontinuously

Each element has its own,unique spectrum

Chapter 5 16

In 1913, Niels Bohr suggested a new model of the atom that explained why

hydrogen had a discrete line spectrum rather than a continuous spectrum.

Bohr's basic theory: electrons in atoms can only be at certain energy levels,

and they can give off or absorb radiation only when they jump from one level

to another.

The Bohr Model of the Atom

In his model that an atom consists of an

extremely dense nucleus that is surrounded by

electrons that travel in set orbits around the

nucleus.

He hypothesized that the energy possessed by

these electrons and the radius of the orbits arequantized, meaning it is limited to specific

values and is never between those values.

These orbits were of varying energies,

dependent on their distance from the nucleusThe Gobstopper Model


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Chapter 5 17

Quantized versus Continuous

The quantum concept means that the energy of the electron andits radius of orbit around the nucleus is limited to specific values

(and cannot be anywhere in between!).

If these values were continuous, they would be free to have any

value.Energy

E1

E2

E3

hv = E3 E2

hv = E3 E1

hv = E2 E1

Chapter 5 18

Atomic Spectra:

The Balmer-Ryberg Equation In 1885, Johann Balmer determined that the pattern of the atomic

spectra of hydrogen could be predicted by a mathematical formula.

Balmer determined that the wavelength or frequency of the lines inthe spectrum could be expressed by the following equations:

In addition to the lines seen in the visible region (Balmer series),there are additional sets of lines found in the UV region (Lymanseries) and the IR region (the Paschen series).

All conform to the above equations.

(E =RH 1ni21

nf2

- )RHis the Ryberg constant (2.18 x 10

-18 J)

ni and nfare integers with nf> ni


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Chapter 5 19

The Bohr Theory: Problems

Bohrs theory only works forhydrogen atoms.

Once you have more than oneelectron, the calculations for theelectron energy and the orbit radii

breakdown.

Therefore, a new theory needed to bedeveloped for multi-electronelements.

Bohrs theory did make twoimportant contributions:

It suggested a reasonable explanationfor the discrete line spectrum of the

elements It introduced the idea of quantized

electron energy levels (orbits!)

Chapter 5 20

The Quantum Mechanical Model of the Atom

In the 1920s Erwin Schrdinger appliedthe principles of wave mechanics toatoms and developed the Quantum

Mechanical Model of the Atom

Basically, Schrdinger said to give up on theidea of literal orbits for the electrons andinstead concentrate on the electron as awave.

This theory builds on Bohrs idea ofquantized energy levels (orbits) and addsadditional requirements for electronlocation and energy.


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Chapter 5 21

Quantum Mechanics

Werner Heisenberg (19011976): supported this

idea by showing that it is impossible to know (or

measure) precisely both the position and velocity

(or the momentum) at the same time.

The simple act of seeing an electron would change its energy and

therefore its position. Think about pinpointing a fly on the wall. What happens when you try to

swat it?

)()4()(:positionselectron'inyUncertaint

4

))((:PrincipletyUncertainHeisenberg

m

hx

hmx

Chapter 5 22

Working with Heisenbergs Principle, Schrdinger

developed a compromise which calculates both the energy

of an electron and the probability of finding an electron at

any point in the molecule.

This is accomplished by solving the Schrdinger equation,

resulting in the wave function, .

These regions were termed orbitals

Quantum Mechanics

Wave

Equation

Wave Function

or Orbital ()

Probability of finding

an electron in aregion of space (2)

solve


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Chapter 5 23

Quantum Numbers

The wave function contains three variablesknown as the Quantum Numbers whichdescribe the size, energy, shape and position ofthe orbitals.

n is the principal energy level (Bohrs Orbits!)

l is the sublevel

ml is the orbital

These numbers serve as an address of the

probable location of the electron.

Chapter 5 24

Principal Quantum Number (n) The Principal Quantum Number (n)

provides info about the distance ofthe electron from the nucleus As n increases, the number of allowed

orbitals also increases as does the sizeof those orbitals.

This increased size allows the electronto reside further from the nucleus

As the electron moves away from thenucleus its energy increases, therefore nalso indicates the energy of electrons

We often state that electrons andorbitals denoted by the same n valueare in the same shell

Allowed Values: n = 1, 2, 3, ... never 0


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Chapter 5 25

Angular-Momentum Quantum Number (l)

The angular momentumquantum number (l) defines the

three dimensional shape of the

orbitals found within a particular

shell

These discrete sets of orbitals

are calledsublevels

The number of sublevels within

a shell is equal to the principal

quantum number (n)

lValue: 0 1 2 3

Letter Used: s p d f

increasing energy

Allowed Values:

l= 0, 1, 2 ... n - 1

Chapter 5 26

Magnetic Quantum Number (ml) The Magnetic Quantum Number (ml) describes the orientation in 3D

space of the sublevel, thereby denoting a specific orbital.

The number of orientations (and therefore orbitals) per sublevel is

determined by the equation:

2l + 1

2l + 1

3 (f)2 (d)1 (p)0 (s)l

Allowed Values: ml= - l, , + l

x axis

y axis

z axis


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Chapter 5 27

Spin Quantum Number (ms)

ThePauli Exclusion Principle statesthat no two electrons can have the

same four quantum numbers

This results in no more than doubleoccupancy in any one orbital

only two electron per orbital and

they must have opposite spins

Chapter 5 28

Electron Occupancy in Sublevels ThePauli Exclusion Principle states that an orbital can hold

up to two electrons

The maximum number of electrons in each of the sublevels

depends on the number of orbitals within that sublevel:

Thes sublevel holds a maximum of 2 electrons (1 orbital).

Thep sublevel holds a maximum of 6 electrons (3 orbitals).

The dsublevel holds a maximum of 10 electrons (5 orbitals).

Thefsublevel holds a maximum of 14 electrons (7 orbitals)

The maximum electrons per principle quantum level (n) is

obtained by adding the maximum number of electrons in eachsublevel.


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Chapter 5 29

Quantum Number Combinations

Why cant an electron have the following quantum numbers?

(a) n = 2, l= 2, ml= 1 (b) n = 3, l= 0, ml= 3

(c) n = 5, l= 2, ml= 1

Give orbital notations for electrons with the following

quantum numbers:

(a) n = 2, l= 1, ml= 1 (b) n = 4, l= 3, ml= 2

(c) n = 3, l= 2, ml= 1

Chapter 5 30

Shapes of the Orbitals Each orbital has a specific shape determined by its

angular momentum quantum number (l).

As you increase the principle quantum number (n), theorbitals increase in size but not shape!

Remember, these orbitals represent a region in spacewhere there is a high probability of finding the electron.

These are not discrete locations!

They are also not pictures of the path that an electron followsaround a nucleus


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Chapter 5 31

l= 0 : The s Orbitals

The s orbitals are spherical, meaning the probability of finding the

electron depends only on distance from the nucleus, not direction.

The value of2 is greatest near the nucleus then drops off as you

move away.

It never reaches zero however so there is technically no definite

boundary to the atom.

Chapter 5 32

l= 1 : The p Orbitals The p orbitals are dumbbell shaped with their electron density concentrated in

identical lobes residing on opposite sides of a nodal plane.

This shape means that a p electron will never be found near the nucleus.

The two lobes of a p orbital have different phases (are opposite in sign) which

becomes important in bonding among atoms.

The three orientations (ml= -1, 0, +1) are 90 differentials along the x, y and

z axes.

The orbitals are designated px (along the x axis), py (along the y axis), and pz(along the z axis)


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Chapter 5 33

l= 2 : The d Orbitals

Chapter 5 34

Effective Nuclear Charge (Zeff) The nuclear charge (Z) of an atom is

determined by the number of protons foundin the nucleus

It is felt by the electrons as an attraction

Multiple electrons in an atom lead to ashielding effect on the outer electrons.

This electron shielding (S) leads to energy

differences among orbitals within a shell.

Net nuclear charge felt by an electron is

called the effective nuclear charge (Zeff).

Zeff= Z + S


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Chapter 5 35

Effective Nuclear Charge (Zeff)

Note that the 3d

sublevel is actually

higher in energy than

the 4s sublevel.

WHY??

Chapter 5 36

The Modern Periodic Table

H.G.J. Moseley discovered that the nuclear charge increased by 1 for each element in theMendeleevs table.

He concluded that the changingatomic number rather than the changing mass explained therepeating trends of the elements

Theperiodic law states that the properties of elements recur in a repeating pattern when arrangedaccording to increasing atomic number.

With the introduction of the concept of electron energy levels byNiels Bohr, the periodic table tookits current arrangement.

http://www.webelements.com/


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Chapter 5 37

Electron Configurations

Many of an elements chemical

properties depend on its electron

configuration

The electron configuration of an atom is a shorthand method ofwriting the location of electrons by sublevel.

The principal quantum level (n) is written first, followed by the

letter designation of the sublevel (l) then a superscript with the

number of electrons in the sublevel.

There are rules for the order and manner that each sublevel isfilled called the Aufbau Principle.

Chapter 5 38

Electron Configurations: Aufbau Principle

Pauli Exclusion Principle:No two electrons in an atom can

have the same quantum numbers (n, l, ml, ms).

Hunds Rule: When filling orbitals in the same sublevel,

maximize the number of parallel spins (so fill then pair!).

Rules of Aufbau Principle:

1. Lower n orbitals fill first.

2. Each orbital holds two electrons; each with different ms.

3. Half-fill degenerate orbitals before pairing electrons.


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Chapter 5 39

Using the Periodic Table for Electron Configurations

The periodic table took on its current shape once the quantum model of the

atom was developed.

You can use it to fill up your sublevels and orbitals to build your electron

configurations.

Chapter 5 40

Writing Electron Configurations Step 1: Locate the element on the periodic table.

Step 2: Determine the number of electrons the element

has: Iron has 26 electrons

Step 3: Starting at the beginning of the Periodic Table,

move left to right across the periods, filling each sublevel

with electrons until you reach the location of your

element:

Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d6

Step 4: Check that the sum of the superscripts equals the

atomic number of iron (26).


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Chapter 5 41

An Alternative Method

1s2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d7s 7p

Increasing

Energy

[He]

[Ne]

[Ar]

[Kr]

[Xe]

[Rn]

Core

Chapter 5 42

Li 1s2 2s1

1s 2s

Be 1s2 2s2

1s 2s

B 1s2 2s2 2p1

1s 2s 2px 2py 2pz

C 1s2 2s2 2p2

1s 2s 2px

2py

2pz

Writing Electron Configurations


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Chapter 5 43


N 1s2 2s2 2p3

1s 2s 2px 2py 2pz

O 1s2 2s2 2p4

1s 2s 2px

2py

2pz

Ne 1s2 2s2 2p5

1s 2s 2px

2py

2pz

S [Ne] [Ne] 3s2 3p4

3s 3px

3py

3pz

Chapter 5 44

Give the ground-state electron configurations for:

Ne (Z= 10) Mn (Z= 25) Zn (Z= 30)

Eu (Z= 63) W (Z= 74) Lw (Z= 103)

Identify elements with ground-state configurations:

1s2

2s2

2p4

1s2

2s2

2p6

3s2

3p6

4s2

3d10

4p6

5s2

4d6

1s2 2s2 2p6 [Ar] 4s2 3d1 [Xe] 6s2 4f14 5d10 6p5



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Chapter 5 45

Exceptions to the Filling Order

When filling the d sublevel, exceptions occur for thechromium (Cr) and copper (Cu) families:

4s 3d 4p

4p4s 3d

4s 3d 4p

4s 3d 4p

Cr

Cu

Chapter 5 46

Periodic Trends The arrangement of the periodic table means that

the physical properties of the elements follow a

regular pattern.

Some trends include:

Atomic Radius (end of Chapter 5)

Ionization Energy (Chapter 6) Electron Affinity (Chapter 6)

Electronegativity (Chapter 7)


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Chapter 5 47

Atomic Radius

An atomsatomic radius is the distance from the nucleus to theoutermost electrons.

Why do you think the radiusincreases in this way?

chapter 5 periodicity and atomic structure

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