Models of the Atom Chapter 5: Electrons in...

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Chapter 5: Electrons in Atoms


What is Light? Electromagnetic Radiation is a fancy shmancy sciency

way of saying Light

Light is a form of Energy. It is NOT made of matter.

It can act like a wave (like ripples in water) or like a

particle

Waves have 3 properties:

Wavelength (): length of one wave from crest to crest, trough to

trough or node to node (m or nm)

Frequency (): number of waves per second (waves/sec or Hz)

Speed (c): light travels at the same speed all the time:

○ 3.00 x 108 m/s


a) What would be the of this wave?

b) What would happen to the if you

shortened the ?

c) What would happen to the Energy of

this wave is you shortened the ?

0 sec 2 sec


How is related to ?

Shorter wavelengths =

HIGHER FREQUENCY

Longer wavelengths =

LOWER FREQUENCY


Wave model of light

According to the wave model of light, light

consists of electromagnetic waves.

Electromagnetic radiation (EMR) includes

Radio waves, Microwaves, Infrared waves, Visible

light, Ultraviolet waves, X-rays, and Gamma rays.

All electromagnetic waves travel in a vacuum at

a speed of 3.00 108 m/s (the speed of light)


The Electromagnetic Spectrum


The Electromagnetic Spectrum

Most light is not visible.

The frequencies of light waves range from

3 x 106 waves/sec to 3 x 1022 waves/sec

any wave that has less or more waves/sec than this is NOT light.

The wavelengths of light waves range from

102 to 10-14 meters long

any wave that is longer or shorter than this range is NOT light.

Notice: Visible light is only a small portion of the EMS. It ranges from 700 nm- 380 nm ROYGBIV


The Visible Spectrum

White light (like sunlight) is actually made of ALL the

visible wavelengths mixed together.

If passed through a prism, then the different colors

can be separated into a spectrum!!!


DOUBLE RAINBOW ALL THE WAY!


Calculating type of EMR

The product of the frequency and wavelength always

equals a constant (C), the speed of light.

The speed of light equals 3.00x108 m/s

Ex: What is the wavelength when the frequency of

the wave is 2.4x1010 Hz? What type of radiation is

this?


Sample Problems:

What is the frequency for a wavelength of 5.00 x

10-8 m? In what range of the spectrum is it?

What is the wavelength and type of radiation for a

frequency of 3.83 x 1010 s-1?


Atomic Spectra A prism will separate any light into the colors it

contains. When white light passes through a prism, it

produces a spectrum of colors that when mixed, make

up that light.


Atomic Spectra When light from a helium lamp passes

through a prism, only specific frequencies

of light are produced.


All elements, when “burned” will emit a

spectrum that is unique to that element.


WHY?


Remember The Bohr Model?

That electrons were

found in certain circular

orbits around the

nucleus

Called the Planetary

Model

5.1


Bohr’s Explanation of Atomic

Spectra cont…

In the Bohr model, the electrons in an atom can have

only certain specific energies.

The specific energy an electron has determines where

it orbits around the nucleus, or what ENERGY LEVEL it

is orbiting in.

n = energy level


Bohr’s Explanation of Atomic Spectra

cont… When the electron has its lowest possible energy, the

electron is said to be in its ground state

Any state that is above the ground state in energy is called

an excited state Absorbing energy raises the electron from the ground state to an excited state.

Photon

released

Energy

Absorbed

Ground state

Excited state



Spectra cont…

The energy absorbed has to be the exact right amount to move the electron from one level to another.

○A quantum of energy is the amount of energy required to move an electron from one energy level to another energy level.


• A quantum of energy in the form of light is then emitted when the electron drops back to a lower energy level.

• An atom’s radius increases

when absorbing and decreases when emitting why?


Spectra cont…


Atomic Spectra

• Bohr’s Model explained

why Hydrogen’s atomic

spectrum gives off three

lines.

• Why?


Atomic Spectra

When an atom absorbs energy:

the electrons in the outer part of the atom get “excited” and move into higher energy levels.

These electrons then lose that energy by emitting light when they return to the lower energy level they once occupied.

The frequencies of light emitted by an element separate into discrete lines to give the atomic emission spectrum of the element.

http://www.visionlearning.com/library/module_viewer.php?mid=51


Atomic Spectra

The atomic emission spectrum for each element are unique to that element. Like a fingerprint.

They are used to identify elements.

The composition of the universe comes from the study of the emission spectra of stars!


Particle Nature of light

In 1900 Max Planck determines that:

Light is NOT a continuous beam of energy

It is a beam of zillions of “Packets” of Energy

○ One of these packets of energy is called a PHOTON

○ These photons have a “particle nature”, meaning

they sometimes ACT LIKE a particle (have collisions

with other particles), even though they are not made

of matter.

http://ethel.as.arizona.edu/~collins/astro/subjects/electromag5.html


Particle Nature of Light cont… In 1905 Einstein confirmed Planck’s idea of the

particle nature of light with the Photoelectric Effect. When u.v. light (photons) is shined on certain sheets of metal,

electrons are “knocked” out of the atoms in the metal.

○ A stream of electrons is an electrical current.

Increasing the frequency (ν) of the incoming light (photons) causes electrons to eject from the metal faster, increasing the energy of these electrons.


Energy of a Photon

The Energy of a photon can be calculated:

• E is the energy of the photon (in Joules or J)

• h is Planck’s constant and is equal to 6.63 x 10-34 J-s

• ν is the frequency of light (w/s or Hz or S-1)

What happens to the energy of a photon if you increase

the frequency?

E = hν


Wave Particle Duality

Einstein confirmed that Light has a

DUAL NATURE:

Particle nature:

Photons have

collisions with

other particles

(electrons)

Wave nature:

Photons have a

frequency,

wavelength, and

amplitude.

Light acts as both particle and wave!!!!


The Development NEW Atomic Models

Based on this new knowledge of how light acts, new

models of the atom were developed.

Why was Bohr’s model considered inadequate?

It couldn’t explain why when different elements,

when energized, glowed different colors and gave

off different atomic spectra.

5.1


The Bohr Model

Like the rungs of this

strange ladder, the energy

levels in an atom are not

equally spaced.

The higher the energy

level occupied by an

electron, the less energy it

takes to move from that

energy level to the next

higher energy level.

Why?

5.1


Color of Photons Emitted

Lowest energy photons are red because they have the lowest frequency

Highest energy photons are violet because they have the highest frequency

Energy/frequency Increases


The Quantum Mechanical Model

Bohr’s model failed to explain the energies absorbed

and emitted by atoms with more than one electron.

So, The Quantum Mechanical Model was developed.

5.1



Like Bohr’s model the electrons are in certain Energy Levels

But instead of orbiting in fixed paths, they swarm in certain regions (Electron clouds) within those energy levels.

5.1



The particular location of an electron within those energy levels is based on probability.

ex: The propeller blade has the same probability of being anywhere in the blurry region, but you cannot tell its location at any instant. The electron cloud of an atom can be compared to a spinning airplane propeller.

5.1



Electrons have a probability of being found in some regions more than others. This can be represented by an electron cloud that is more dense in some places, and less dense in others.

The cloud is more dense where the probability of finding the electron is high.

5.1


Images of real atoms


Organization of e- in atoms

There are SEVEN Principle ENERGY

(Quantum) LEVELS (n)

Each Energy Level is made up of

SUBLEVELS

The number of sublevels an energy level

contains is equal to the principle quantum

number of that level. ○Ex: the 1st Energy Level has 1 sublevel

the 2nd Energy Level has 2 sublevels etc…

5.1


Sublevels and Orbitals

○The types of sublevels are: 1st sublevel: s

2nd sublevel: p

3rd sublevel: d

4th sublevel: f

○Sublevels are made up of ORBITALS An orbital is the region of space in which there is a high probability of finding an electron.

Each orbital can hold 2 electrons

Each type of sublevel has different numbers of orbitals

Orbitals have certain shapes and orientations and represent an area where electrons are likely to be found.


Atomic Orbitals

s sublevels contain 1 orbital

How many electrons in each? How many total?

p sublevels contain 3 orbitals


d sublevels contain 5 orbitals


f sublevels contain 7 orbitals



http://www.youtube.com/watch?v=VfBcfYR1VQo&fe

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http://www.youtube.com/watch?v=sMt5Dcex0kg&fea

ture=related

http://www.youtube.com/watch?v=VfBcfYR1VQo&feature=related

http://www.youtube.com/watch?v=VfBcfYR1VQo&feature=related


Summary of sublevels

Orbital and Electron Capacity for the First Four Sublevels

Sublevel # of orbitals Maximum number of

electrons

s 1 2

p 3 6

d 5 10

f 7 14

Models of the Atom > Summary of e- in Energy

Level Summary of Atomic Energy Levels, Sublevels and Orbitals

Principle

Energy

Levels (n)

#of

sublevels

Type of

sublevel

#of Orbitals

per sublevel

# of

electrons per

sublevel

Total # of

electrons per

energy level

1 1 1s 1 2 2

2 2 2s 1 2 8

2p 3 6

3 3 3s 1 2

18 3p 3 6

3d 5 10

4 4 4s 1 2

32 4p 3 6

4d 5 10

4f 7 14


Atomic Orbitals ○ Different atomic orbitals are denoted by letters.

The s orbitals are spherical, and p and d orbitals

are dumbbell-shaped.

5.1


Examples of Orbitals

http://daugerresearch.com/orbital

s/


Electron Configuration

Electron configuration is the arrangement

of electrons in an atom

There are 3 rules for Electron

Configuration…


Rule #1(Aufbau principle)

Electrons enter the lowest available

energy orbital first


Rule # 2 (Pauli exclusion principle)

There are a maximum of 2 electrons per orbital which means if there are two electrons in the same orbital they must have opposite spins

One spins clockwise, the other counter clockwise


Rule #3 (Hund’s Rule)

For p, d, f orbitals 1 e- enters each orbital

until all contain 1 e-

Then a 2nd e- added

1 at a time


Example: Nitrogen atomic # 7

1s

2s

2p 2p 2p

Or Shorthand:

1s2 2s2 2p3


Electron configuration shorthand

1s2

Principle

energy

level (n)

Sublevel

(s,p,d,f)

Number of

electrons

Models of the Atom > Electron configuration using your

Periodic Table

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