Electronic Structureand the Periodic Table

Unit 6 Honors Chemistry

Electromagnetic Waves:

Electromagnetic waves:progressive, repeating disturbances that come from the movement of electric charges

Electromagnetic Waves & Light

Wavelength and FrequencyWavelength and Frequency

Wavelength (, lambda): distance between any two points in a wave

measured in any distance unit

(mainly nm or m:

1 nm = 1x10-9 m)

Wavelength Can be Measured in One of Two Ways…

Wavelength and FrequencyWavelength and Frequency

Frequency (; pronounced nu): the number of cycles of the wave that pass through a

point in a unit of timeMeasured in sec-1 (/sec)

1 sec-1 = 1 Hertz (Hz)

Illustration of Frequency

Wavelength is indirectly proportional to frequency

As Wavelength increases, frequency _________________.

As Wavelength decreases, frequency _________________.

AmplitudeAmplitude

Note: height of wave is amplitude (intensity or brightness of wave)

Amplitude is INDEPENDENT of frequency or wavelength!

SpeedSpeed

Speed (c): The speed of light!

c = 3.00 x 108 m/s

(rounded to 3 sig figs)

Equation Equation

One equation relates speed, frequency and wavelength:

c =

Example The wavelength of the radiation which

produced the yellow color of sodium vapor light is 589.0 nm. What is the frequency of this radiation?

The electromagnetic spectrum complete range of wavelengths and frequencies mostly invisible

What is color?

TED Talk: What is color?

The visible/continuous spectrum

continuous spectrum: components of white light split into its colors, ROY G BIV from 390 nm (violet) to 760 nm (red)

Line Spectra

Pattern of lines produced by light emitted by excited atoms of an elementunique for every elementused to identify unknown elements

How do we see color?

TED Talk: How we see color

Max Planck

Light is generated as a stream of particles called PHOTONS

Equation: E (Energy of a photon)= h

(h =Plank’s constant= 6.626x10-34Js)

Low frequency, high Low frequency, high λλ, low E., low E.Low frequency, high Low frequency, high λλ, low E., low E.

High frequency, low High frequency, low λλ, high E., high E.High frequency, low High frequency, low λλ, high E., high E.

E = h • E = h • E = h • E = h •

Relationships in Planck’s Eqn.Relationships in Planck’s Eqn.Relationships in Planck’s Eqn.Relationships in Planck’s Eqn.

Photoelectric effect – Nobel Prize in Physics 1921 to Einstein

Occurs when light strikes the surface of a metal and electrons are ejected.

Practical uses:

Automatic

door openers

Photoelectric Effect: ConclusionPhotoelectric Effect: Conclusion

Light not only has Light not only has wave wave propertiesproperties but also has but also has particle particle

properties. These massless properties. These massless particles, called particles, called photonsphotons, , are are

packets of energy.packets of energy.

Example 6.2

Using the frequency calculated in the previous example, calculate the energy, in joules, of a photon emitted by an excited sodium atom. Calculate the energy, in kilojoules, of a mole of excited sodium atoms.

Bohr’s Hydrogen Atom: Bohr’s Hydrogen Atom: A Planetary Model A Planetary Model

Niels Bohr: ProposedNiels Bohr: Proposedplanetary model.planetary model.

Electrons “orbit” the nucleus like planets Electrons “orbit” the nucleus like planets around the sun.around the sun.

NOT current model of atom but used to NOT current model of atom but used to explain some features of atom.explain some features of atom.

Ground State vs. Excited State

ground state: all electrons in lowest possible energy levels

excited state: an electron that has absorbed energy and moved to a higher energy levelThis is a temporary state!!

Explanation of Line Spectra & Equation

Niels Bohr

Energy of an electron is quantized: can only have specific values.

Energy proportional to energy level.

Explanation of Line Spectra

Electron will drop from excited state to ground state and will emit energy as a photon.

Explanation of Line Spectra

Type of photon emitted by electron depends on energy difference of energy levels

Elevel = -RH 1 – 1 (nhi)2 (nlow)2

AND Elevel = h = hc/

(h: Planck’s constant, 6.626 x 10-34 J sec/photon)

Flaw in Bohr’s Model

Only works well for 1 electron species (H atom).

Does not explain fine structure of line spectra.

Wave-Particle Duality

Light has properties of both WAVES and PARTICLES.

most matter has undetectable wavelengths (1000 kg car at 100 km/hr has = 2.39 x 10-38 m)

This work led to the development of the electron microscope

Quantum Mechanics

Quantum mechanics:

atomic structure based on wave-like properties of the electron

Schrödinger: wave equation that describes hydrogen atom

Heisenberg Uncertainty Principle

The exact location of an electron cannot be determined (if we try to observe it, we interfere with the particle)

You can know either the location or the velocity but not both

Electrons exist in electron clouds

and not on specific rings or orbits

Quantum Numbers Four quantum numbers are a mathematical

way to represent the most probable location of an electron in an atom

analogy...state = energy level, n

city = sublevel, l

address = orbital, ml

house number = spin, ms

Principal Quantum Number: n

Always a positive integer (1,2, 3…7)

Indicates size of orbital, or how far electron is from nucleus

Similar to Bohr’s energy levels or shells

Larger n value = larger orbital or distance from nucleus

The Periodic Table and The Periodic Table and ShellsShells

The Periodic Table and The Periodic Table and ShellsShells

n = 1n = 1

n = 2n = 2

n = 3n = 3

n = 4n = 4

n = 5n = 5n = 6n = 6n = 7n = 7

n = row number on periodic table for a given element

Angular Momentum Quantum Number: l

positive integer from zero to n-1 Sublevel within an energy level; indicates

shape of orbital 0 = s 1 = p 2 = d 3 = f

Types of Types of SublevelsSublevels

s s p p d d

Magnetic Quantum Numbers: ml

integer from -l to +l Indicates orientation of orbital in space

Orbital = electron containing area

Spin Quantum Number: ms

Two values only: + ½ or -½ 2 electrons max. allowed in each orbital (Pauli Exclusion Principle)

Indicates spin of electron; spins of each electron must be opposite

n ---> level1, 2, 3, 4, ...

l ---> sublevel 0, 1, 2, ... n - 1

ml ---> orbital -l ... 0 ... +l

ms ---> electron spin +½ and -½

n ---> level1, 2, 3, 4, ...

l ---> sublevel 0, 1, 2, ... n - 1

ml ---> orbital -l ... 0 ... +l

ms ---> electron spin +½ and -½

REVIEW: REVIEW: QUANTUM NUMBERSQUANTUM NUMBERSEvery Electron has four!Every Electron has four!

OrbitalsOrbitalsNo more than 2 e- assigned to an orbitalNo more than 2 e- assigned to an orbitalOrbitals grouped in s, p, d (and f) Orbitals grouped in s, p, d (and f)

subshellssubshells

s orbitalss orbitals

d orbitalsd orbitals

p orbitalsp orbitals

Capacities of levels, sublevels, and orbitals—see packet

Example

Example 6.6 Give the n and l values for the following orbitals:

a. 3p

b. 4s

Example

Example 6.8 What are the possible ml values for the following orbitals:

a. 3p

b. 4f

Shapes of Atomic Orbitals

Shapes of Atomic Orbitals

s = spherical

p = peanut

d = dumbbell (clover)

f = flower

Multielectron Atoms

In the hydrogen atom the subshells (sublevels) of a principal energy level or shell are at the same energy level.

Previous Equation: En = –RH /n2

Multielectron Atoms

In a multielectron atom, only the orbitals are at the same energy level: the sublevels are at different energy levels!

The increasing energy order of sublevels is generally:

s < p < d < f

Overlapping subshells

At higher energy levels, sublevels overlap.

Note:

4s vs. 3d!

Introduction to Electron Configuration

Definition: describes the distribution of electrons among the various orbitals in the atom

Represents the most probable location of the electron!

EOS

Electron Configurations

The system of numbers and letters that designates the location of the electrons

3 major methods: Full electron configurations Abbreviated/Noble Gas configurations Orbital diagram configurations

Full or Complete Electron Configuration (uses spdf)

Uses numbers to designate a principal energy level and the letters to identify a sublevel; a superscript number indicates the number of electrons in a designated sublevel.

EOS

Rules for Electron Configurations

The Aufbau principle:Electrons fill from the lowest energy level to the highest (they don’t skip around)

1s22s22p63s23p64s23d10etc.

Pauli Exclusion Principle

No two electrons in the same atom can have the same set of 4 quantum numbers.

That is, each electron has a unique “address”

In other words, the maximum # of electrons an orbital can hold is 2 e- (one with ms = +1/2 and one

with ms = -1/2)

HUND’S RULE

Orbitals of equal energy in a sublevel must all have 1 electron before the

electrons start pairing up

a.k.a “creepy person on the bus rule”

*** also electrons in half-filled orbitals have same spin

Why are these incorrect?

Why are these incorrect?

Why are these incorrect?

Full Electron Configuration

Example Notation: 1s2 2s1 (Pronounced “one-s-two, two-s-one”)

A. What does the coefficient mean?Principle energy level

B. What does the letter mean?Type of orbital (sublevel)

C. What does the exponent mean?# of electrons in that orbital

Steps to Writing Full Electron Configurations

1. Determine the total number of electrons the atom has (for neutral atoms it is equal to the atomic number for the element).

Example: F atomic # = # of p+ = # of e- =

2. Fill orbitals in order of increasing energy (see Aufbau Chart).

3. Make sure the total number of electrons in the electron configuration equals the atomic number.

Aufbau Chart (Order of Energy Levels)

When writing electron configurations:

d sublevels are n – 1 from the row they appear in

f sublevels are n – 2 from the row they appear in

Writing Electron Configurations

Writing Electron Configurations

Nitrogen:

Helium:

Phosphorous:

Rhodium:

Bromine:

Cerium:

Abbreviated/Noble Gas Configuration

i. Where are the noble gases on the periodic table?

ii. Why are the noble gases special?

iii. How can we use noble gases to shorten regular electron configurations?

Abbreviated/Noble Gas Configuration

Example: Barium

1.Look at the periodic table and find the noble gas in the row above where the element is.

2.Start the configuration with the symbol for that noble gas in brackets, followed by the rest of the electron configuration.

Abbreviated/Noble Gas Configuration

Practice! Write Noble Gas Configurations for the following elements:

Rubidium:

Bismuth:

Arsenic:

Zirconium:

Writing Electron Configurations

Writing Electron Configurations

Another way of writing

configurations is called an orbital

diagram.

(also called orbital notation or orbital

diagrams)

Another way of writing

configurations is called an orbital

diagram.

(also called orbital notation or orbital

diagrams)

Arrowsdepictelectronspin

ORBITAL BOX NOTATIONfor He, atomic number = 2

1s

21 s

Arrowsdepictelectronspin

ORBITAL BOX NOTATIONfor He, atomic number = 2

1s

21 s

One electron has n = 1, One electron has n = 1, ll = 0, m = 0, mll = 0, = 0, mmss = + = + ½½

Other electron has n = 1, Other electron has n = 1, ll = 0, m = 0, mll = 0, = 0, mmss = - = - ½½

Orbital Diagrams

Orbital diagrams use boxes (sometimes circles) to represent energy levels and orbitals. Arrows

are used to represent the electrons.

= orbital

sublevels

Orbital Diagrams

Don’t forget - orbitals have a capacity of two electrons!! Two electrons in the same orbital must have opposite

spin so draw the arrows pointing in opposite directions.

Example: oxygen 1s22s22p4

1s

2s

2p

Incr

easi

ng E

nerg

y

Drawing Orbital Diagrams

1. First, determine the electron configuration for the element.

2. Next draw boxes for each of the orbitals present in the electron configuration.

Boxes should be drawn in order of increasing energy (see the Aufbau chart).

3. Arrows are drawn in the boxes starting from the lowest energy sublevel and working up. This is known as the Aufbau principle.

Add electrons one at a time to each orbital in a sublevel before pairing them up (Hund’s rule)

The first arrow in an orbital should point up; the second arrow should point down (Pauli exclusion principle)

4. Double check your work to make sure the number of arrows in your diagram is equal to the total number of electrons in the atom.

# of electrons = atomic number for an atom

Electron Configurations for Nitrogen

Electron Configurations for Nitrogen

Electron Configurations for Nickel

Electron Configurations for Nickel

LithiumLithium

Group 1A

Atomic number = 3

1s22s1 ---> 3 total electrons

1s

2s

3s3p

2p

BerylliumBeryllium

Group 2A

Atomic number = 4

1s22s2 ---> 4 total electrons

1s

2s

3s3p

2p

BoronBoron

Group 3A

Atomic number = 5

1s2 2s2 2p1 --->

5 total electrons

1s

2s

3s3p

2p

CarbonCarbon

Group 4A

Atomic number = 6

1s2 2s2 2p2 --->

6 total electrons

Here we see for the first time Here we see for the first time

HUND’S RULEHUND’S RULE. .

1s

2s

3s3p

2p

NitrogenNitrogen

Group 5A

Atomic number = 7

1s2 2s2 2p3 --->

7 total electrons

1s

2s

3s3p

2p

OxygenOxygen

Group 6A

Atomic number = 8

1s2 2s2 2p4 --->

8 total electrons

1s

2s

3s3p

2p

FluorineFluorine

Group 7A

Atomic number = 9

1s2 2s2 2p5 --->

9 total electrons

1s

2s

3s3p

2p

NeonNeon

Group 8A

Atomic number = 10

1s2 2s2 2p6 --->

10 total electrons

1s

2s

3s3p

2p

Note that we Note that we have reached the have reached the end of the 2nd end of the 2nd period, and the period, and the 2nd shell is full!2nd shell is full!

Exceptions to the Filling Order Rule (Cr, Cu)—these will not be on test!

Valence electrons

Importance and definition:

Definition: Electrons in the outermost energy levels; they determine the chemical properties of an element.

Write the noble gas configuration...the valence electrons are the ones beyond the core.

Example: Sulfur

Valence Electrons and Core Configuration (Shorthand)

What is the shorthand notation for S?

EOS

Sulfur has six valence electrons

Configurations of Ions

Cations: Formed when metals lose e– in highest principal energy level.

Example:

(Z = 11) Na

EOS

(Z = 11) Na+

Configurations of Ions

Anions: Formed when non-metals gain e– to complete the p sublevel.

Example:

EOS

-Z= 18 Cl-

Transition Metals

Transition metals (and p block metals) lose e– from the highest principal energy level

(n) FIRST, then lose their d electrons!

EOS

Zr: [Kr] 5s24d2

Zr+2 : [Kr] 4d2

Isoelectronic Species

Definition: Ions or atoms that have the same number of electrons

Example: Neon, O2-, F-, Na+, Mg2+, Al3+

all have the same configuration (1s22s22p6) and are isoelectronic

Electron Spin and Magnetism

•DiamagneticDiamagnetic: NOT : NOT

attracted to a magnetic attracted to a magnetic

fieldfield

•ParamagneticParamagnetic: :

substance is attracted to substance is attracted to

a magnetic field. a magnetic field. •Substances with Substances with

unpaired electronsunpaired electrons are are

paramagnetic.paramagnetic.

Examples

Mg

Cl

Write orbital notation: if it has an unpaired e- it is paramagnetic.

Periodic Properties & Trends

Electronegativity Ability of an atom to pull e- towards itself

Increases going up and to the right Across a period more protons in nucleus = more

positive charge to pull electrons closer Down a group more electrons to hold onto =

element can’t pull e- as closely

Periodic Properties & Trends

Electronegativity Ability of an atom to pull e- towards itself Across a period more protons in nucleus = more

positive charge to pull electrons closer Down a group more electrons to hold onto = protons

in nucleus can’t pull e- as closely

Atomic Radius

Definition:

½ experimental distance between centers of two bonded atoms

Atomic Radius

Trend in a family:

Size increases

down a group.

(More principal

energy levels)

Atomic Radius

Trend in a period:

Size decreases across a period, e- more strongly attracted to nucleus.

Atomic Radius

Transition metals:

Size stays relatively constant across a period; e-

added to inner energy level.

Memory DeviceMemory Device

LLLL: Lower Left, Larger AtomsLLLL: Lower Left, Larger Atoms

Sizes of IonsSizes of Ions

CATIONS are SMALLER than the atoms from which they are formed.

Size decreases due to increasing he electron/proton attraction.

Li,152 pm3e and 3p

Li+, 78 pm2e and 3 p

+

Sizes of IonsSizes of Ions

ANIONS are LARGER than the atoms from which they are formed.

Size increases due to more electrons in shell.

F, 71 pm9e and 9p

F-, 133 pm10 e and 9 p

-

Trends in Ion SizesTrends in Ion Sizes

Active Figure 8.15Active Figure 8.15

Trends in ion sizes are the same Trends in ion sizes are the same as atom sizes.as atom sizes.

First Ionization EnergyFirst Ionization Energy

Definition: energy required to remove an electron from an atom in the gas phase.

Mg (g) + 738 kJ ---> MgMg (g) + 738 kJ ---> Mg++ (g) + e- (g) + e-

First Ionization Energies

EOS

Trend in a group:

Decreases going down a group (e- further away; easier to remove)

Trend in a period:

Increases going across a period (e- held more tightly).

Memory DeviceMemory Device

LLLL: Lower Left, Larger Atoms;Looser

electrons

LLLL: Lower Left, Larger Atoms;Looser

electrons

Mg (g) + 738 kJ ---> Mg+ (g) + e-

MgMg+ + (g) + 1451 kJ ---> Mg(g) + 1451 kJ ---> Mg2+2+ (g) + e- (g) + e-

Definition: energy required to remove 2nd electron from an atom in the gas phase. Takes more energy because e- is removed from increasingly positive ion.

Second Ionization EnergySecond Ionization Energy

Electron Affinity

Some elements GAIN electrons to form anions.

Electron affinity is the energy involved when an atom gains an electron to form an anion.

A(g) + e- ---> A-(g) E.A. = ∆E

Trends in Electron Affinity

Trend in a group:

Affinity for e- decreases

going down a group

Trend in a series or period:

Affinity for e- increases going across a period

Electron Affinity

Note that the trend for E.A.

is the SAME as for I.E.!

Trends in Metallic Properties

Most metallic means easiest loss of electrons!Metals are on left, nonmetals on right of p.t.

A Summary of Periodic Trends

Remember LLLL!!