Modern Atomic Model

Electron modeling…

• To understand electrons, scientists began comparing them to light.

Behavior of light

•Light is a wave – similar to water waves

•Visible light belongs to electromagnetic spectrum

Radiowaves

Microwaves

Infrared .

Ultra-violet

X-Rays

GammaRays

Low energy

High energy

Low Frequency

High Frequency

Long Wavelength

Short WavelengthVisible Light

Spectrum

Behavior of light

• All forms of EMR travel at a constant speed of 3.0 x 108 m/s, when in a vacuum. This value also works for the speed of light in air (because air is mostly a vacuum).

Behavior of lightWaves can be described in terms

of the following:•Wavelength•Frequency•Amplitude•Speed

Behavior of lightWaves can be described in terms of the following:

•Wavelength - , distance between successive crests (or any 2 corresponding points), for visible light = 400-750 nm

• Frequency - the number of waves that pass a given point per second, units are cycles/sec or hertz (Hz), abbreviated - the Greek letter nu

c =

Behavior of light

Behavior of light•Amplitude – height from origin to crest

Behavior of light•Speed – measured in m/s, light moves @ constant speed of 3.0 x 108 m/s, abbreviated as c

Parts of a wave

Wavelength

AmplitudeOrigin

Crest

Trough

Parts of Wave• Origin - the base line of the energy.• Crest - high point on a wave• Trough - Low point on a wave• Amplitude - distance from origin to crest• Wavelength - distance from crest to crest,

abbreviated Greek letter lambda)

Behavior of light

• Because light has a constant speed, we get a relationship between & (c = )

Frequency and wavelength• Are inversely related

– As one goes up, the other goes down.• Different frequencies of light go with

different colors of light.• There is a wide variety of frequencies

– The whole range is called a continuous spectrum

Behavior of lightQuestion 1: If light has = 633 nm,

what is ?

Question 2: Red light travels at 3.0 x 108 m/s and has a of 700 nm. What is ?

Question 3: Violet light has a of 7.5 x 1014 Hz. What is ?

Wave model problemsThe wave model of light worked

well until the beginning of the 20th century. This is because some scientists were observing light and found that what they saw did not fit the wave model.

Wave model problemsBlack body radiation• In 1900, Max Planck was studying

radiation given off when matter was heated. The physics he knew said that matter could absorb or emit any quantity of energy. The results of his experiments did not fit with that idea.

Light is a Particle• Energy is quantized.

– Light is energy– Light must be quantized– These smallest pieces of light are called

photons.• Energy and frequency are directly related.

Wave model problems

•A quantum of light was later called a photon. Radiation is emitted or absorbed in whole numbers of photons.

Wave model problems• To relate the quantum of

energy and the frequency of the radiation, he created the relationship E = h.

Energy and frequency• E = h x

– E is the energy of the photon– is the frequency– h is Planck’s constant

• h = 6.626 x 10 -34 Joules × seconds

Wave model problems• What energy is given off when

your stove coils turn red? (Remember that red light has a frequency of 4.29 x 1014 Hz.)

• Which has greater energy – red or violet light?

Wave model problems

• Planck’s ideas were not immediately accepted. It was not until some time later that Albert Einstein used Planck’s equation to work on solving the photoelectric effect.

Wave model problems Photoelectric effect• Light shining on certain metals

can eject electrons.

Wave model problems Photoelectric effect• The fact that light was able to

knock electrons loose wasn’t a problem. What wave theory couldn’t explain was why only certain frequencies of light (or higher) could knock out electrons.

• Photoelectric Effect Simulation

Wave model problems Photoelectric effect• Einstein proposed that light

consisted of energy quanta that behaved as particles – not waves. The quanta were called photons.

Wave model problems Photoelectric effect• The photoelectric effect problem

was then solved by the idea that radiation is emitted or absorbed in whole numbers of photons or radiation particles.

Wave model problems Photoelectric effect• It was later proven that light

could definitely act as a particle. So, we now have light acting as both a wave and as particles. (This will be the basis for understanding how e- behave.)

Atomic Spectrum

What color tells us about atoms

Prism• White light is made up

of all the colors of the visible spectrum.

• Passing it through a prism separates it.

If the light entering the prism is not white…

• By heating a gas or using electricity, we can get the gas to give off colors

• Passing this light through a prism does something different than white light

Atomic Spectrum• Each element gives off

its own characteristic colors

• Can be used to identify the element

• How we know what stars are made of

Wave model problems Bright line spectrum• Scientists noticed that you could

vaporize an element in a flame to produce different flame colors. You can then use a prism to sort the colors to produce a line spectrum (only certain colors are produced).

Wave model problems Bright line spectrum• Problem: Each element produced

a different line spectrum.

• These are called line spectra

• They are unique to each element.

• These are emission spectra (the light is emitted or given off)

An explanation of Atomic Spectra

Rutherford’s Model• Discovered the

nucleus• Small dense and

positive• Electrons around

nucleus in electron cloud

Bohr’s Model

• Why don’t the electrons fall into the nucleus?

• Move like planets around the sun.• In circular orbits at different levels.• Energy separates one level from another.

Bohr’s Model

Nucleus

Electron

Orbit

Energy Levels

Bohr’s ModelFurther away

from the nucleus means more energy.

There is no “in between” energy

Energy is in LevelsIn

crea

sing

ene

rgy

Nucleus

First

Second

Third

Fourth

Fifth

}

Bohr Model• Niels Bohr was able to explain the

bright line spectrum. To do so, he created a model with the following parts:

•The e- could orbit the nucleus only in allowed paths or orbits.

Bohr Model

•The H atom has set energy possibilities that depend on which orbit the e- occupies.

•The ground state occurs when the e- is in the orbit closest to the nucleus.

Bohr Model•The orbit where the e- is determines the outer dimensions of the atom.

•The energy of the e- increases as it moves into orbits that are farther and farther from the nucleus (excited atom).

Where the electron starts

• The energy level an electron starts from is called its ground state.

Changing the energy• Let’s look at a hydrogen atom

Changing the energy• Heat or electricity or light can move the

electron up energy levels

Changing the energy• As the electron falls back to ground state it

gives the energy back as light

• May fall down in steps– Each with a different energy, frequency, and

wavelength

Changing the energy

The Bohr Ring Atom

n = 3n = 4

n = 2n = 1

{{{

The Bohr Ring Atom

• The farther the electrons fall, the more energy released and higher frequency produced

• All the electrons can move

Bohr Model• Bohr said that the energy of an

electron is quantized (like light) so there have to be energy levels (orbits) where the e- can be. The e- must be given a certain amount of energy to “jump” orbits.

Bohr Model

• If the electron is in the lowest energy level possible (closest to nucleus), that is called the ground state.

Bohr Model

•If energy is put into the e-, it goes to a higher level or an excited state. The spectral lines produced were due to the energy given off when the e- fell.

Bohr Model Problems

• Unfortunately, Bohr’s model only worked for H. What it did do was to get other scientists thinking.

De Broglie•Determined that particles of matter could act as waves.•Described the wavelength of moving particles.

Conclusion:Matter exhibits both wave and

particle properties!

Where are the electrons?• We know they are outside of the

nucleus.• We say they are in an electron cloud.

• But where?

Quantum-mechanical model

Takes into account the following: • Treats e- as waves within the

atom.

Quantum-mechanical model

• e- inside of atoms have specific E and occupy 3-D regions about the nucleus called orbitals. [Orbitals are different than orbits.]

Quantum-mechanical model• The size and shape of the orbitals

depends on the E of the e- that occupy them.

• All orbitals in an atom make up the e- cloud around the nucleus. The e- cloud gives the atom a size and shape.

• The e- can’t be located exactly in the atom. There are areas of probability to find an e-.

The Quantum Mechanical Model• Has energy levels

for electrons.• Contains orbitals of varying shapes and

sizes• It can only tell us the probability of finding

an electron a certain distance from the nucleus.

The Quantum Mechanical Model• The electron is found

inside a blurry “electron cloud”– An area where there

is a chance of finding an electron.

Atomic Orbitals• Principal Quantum Number (n) = the

energy level of the electron.• Within each energy level complex math

describes several shapes.– These are called atomic orbitals

Summary of atomic orbitals

s

p

d

f

# of shapes

Max electrons

Starts at energy level

1 2 1

3 6 2

5 10 3

7 14 4

By Energy Level• First Energy Level• only s orbital• only 2 electrons

• Second Energy Level• s and p orbitals are

available• 2 in s, 6 in p• 8 total electrons

By Energy Level• Third energy level• s, p, and d orbitals• 2 in s, 6 in p, and 10

in d• 18 total electrons

• Fourth energy level• s,p,d, and f orbitals• 2 in s, 6 in p, 10 in d,

and 14 in f• 32 total electrons

By Energy Level• Any thing past the fourth level - not all

the orbitals will fill up.–You simply run out of electrons

• The orbitals do not fill up in a neat order.

Filling order• Lowest energy fill first.• The energy levels overlap• The orbitals do not always fill up order of

energy level.

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p7p

3d

4d

5d

6d

4f

5f

Electron Configurations• The way electrons are arranged in atoms.• Aufbau principle - electrons enter the lowest

energy first– This causes difficulties because of the overlap of

orbitals of different energies.• Pauli Exclusion Principle - at most 2 electrons

per orbital with different spins• Hund’s Rule - When electrons occupy orbitals

of equal energy they don’t pair up until they have to

Electron Configurations

Notations of e- in atoms:• Orbital diagram - unpaired e- represented

by or , paired e- shown as .– Write orbital diagrams for elements 1-10.

• e- configuration notation - no more lines & arrows, uses # of e- in a sublevel as a superscript over the sublevel designation

– Write electron configuration notation for elements 1-10.

Electron Configurations

Notations of e- in atoms:• Shortcut for orbital diagrams &

e- configuration notation– Uses noble gas (group 18

elements) “core”

Noble Gas Shortcuts…

• 1. Find element on periodic table.• 2. Move up 1 row on table and go to Noble

Gas at end of that row• 3. Put this noble gas symbol inside [ ].• 4. Now, write out what is left over after the [ ].

Examples• Li =1s2 2s1

– Noble gas in row above is He– [He] 2s1 is the same as 1s2 2s1

• Be = 1s2 2s2

– Noble gas in row above is He– [He] 2s2 is the same as 1s2 2s2

• Na = 1s2 2s2 2p6 3s1

– Noble gas in row above is Ne– [Ne] 3s1 is the same as 1s2 2s2 2p6 3s1

Electron Configurations

Questions:1) How many sublevels are found in the 3rd energy level? 2) If there were an 8th energy level, how many sublevels would it have?3) How many orbitals are in the s sublevel? p sublevel? d sublevel? f

sublevel?4) If there were a sublevel past the f sublevel, how many orbitals would it

have?5) How many orbitals are in the 2nd energy level?6) How many orbitals are in the 4th energy level?7) How many orbitals would be in the 6th energy level?8) How many electrons are able to go into an orbital?9) How many electrons would there be in the s sublevel? p sublevel? d

sublevel? f sublevel?10) How many electrons would you find in the 1st energy level? 4th energy

level?11) If we had a 5th energy level, how many electrons would it have?

Periodic Table History

History of the Periodic Table

Dobereiner• Triads of elements with shared properties Cannizzaro• method for measuring atomic masses and

interpreting the results of measurementsNewlands• arranged elements by atomic masses,

properties repeated after every 8 elements → law of octaves

History of the Periodic Table

Mendeleev & Meyer• arranged elements according to

the increase in atomic mass

History of the Periodic Table

Mendeleev• left spaces for undiscovered

elements & predicted properties of those elements

• credited with discovering periodicity

History of the Periodic Table

2 Questions:1. Why could most elements be

arranged by increasing atomic mass, but a few could not?

2. What was the reason for chemical periodicity?

History of the Periodic TableMosely• shooting electrons at various metals to

produce X-rays• frequencies of the X-rays were unique to

the metals• assigned a whole number to each

element → atomic numbers» arrange elements by atomic numbers

to get families with similar properties

History of the Periodic Table

Periodic Table• an arrangement of the elements

in order of their atomic numbers so that elements with similar properties fall in the same column (group/family)

History of the Periodic Table

Periodic Law• the physical and chemical

properties of elements are periodic functions of their atomic numbers

Periodic Table Information

Types of Elements

Periodic Table InformationTypes of elements:• Metal – lustrous, good conductors, most are

solids, malleable, ductile• Nonmetal - poor conductors, no luster,

neither malleable nor ductile, most are gases, wide variety of other physical properties

• Metalloid - properties of metals & non-metals

Parts of the Periodic Table

Periodic Table Information

Period• horizontal rows of elementsGroup (family)• vertical columns of elements

Periodic Table InformationGroup (family)• alkali metals – group 1 (except hydrogen)• alkaline earth metals – group 2• transition metals – all of the d-block elements• inner transition metals – all of the f-block elements• metalloids – elements that touch the “staircase”• halogens – group 17• noble gases – group 18 (elements with filled outer

shells of electrons)

Back to Electrons…(electron configurations)

Electron Configuration

Notation(using shortcut)

More Electron ConfigurationsPractice writing shortcut e-

configurations.•Element # 15

•Element # 8

•Element # 34

Group Electron Configuration

More Electron Configurations

Group e- configurations – all elements in family have similar “endings”

Write the shortcut e- configuration for:• Li, Na, K

»all end with __, so group configuration is __

• C, Si, Ge»all end with __, so group

configuration is __

More Electron Configurations

• Element in period 3 with group configuration p4

• Element in period 6 with group configuration s2

• Element in period 2 with group configuration p6

• Element in period 7 with group configuration s1

Practice identifying elements using group configurations:

Periodic Trends

Periodic TrendsElectron configurations are able to

cause periodic variations in elemental properties...

Periodic TrendsValence electrons - electrons that may

be lost/gained/shared when chemical compounds are formed

• Period As we go across a period, the number of valence electrons increases.

• Group As we go down a group, we find that the number of valence electrons stays constant.

Periodic TrendsSize of the atomic radius – one-half the

distance between the nuclei of identical atoms joined in a molecule

• Period As we go across a period, the general trend is for the atomic radii to decrease.

• Group As we go down a group, there is a general increase in atomic radii. This happens because we are seeing more and more energy levels being added.

Periodic TrendsIonization energy – energy required to

overcome nuclear attraction and remove an electron from a gaseous element

• Period As we go across a period, the general trend is for the ionization energy to increase.

• Group As we go down a group, there is a general decrease in ionization energy.

Periodic TrendsElectronegativity - measure of the power

of an atom in a chemical compound to attract electrons

• Period As we go across a period, the general trend is for the electronegativity to increase.

• Group As we go down a group, there is a general decrease in electronegativity.