5.1 Revising the Atomic Model > 1 Chapter 5 Electrons In Atoms 5.1 Revising the Atomic Model 5.2 Electron Arrangement in Atoms 5.3 Atomic Emission Spectra

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Chapter 5Electrons In Atoms

5.1 Revising the Atomic Model

5.2 Electron Arrangement in Atoms

5.3 Atomic Emission Spectra and the Quantum Mechanical Model


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Shown here is a life-sized model of a skier, but not all models are physical. In fact, the current model of the atom is a mathematical model.

CHEMISTRY & YOUCHEMISTRY & YOU

Why do scientists use mathematical models to describe the position of electrons in atoms?


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Understand how the atomic model was revised.

Energy Levels in AtomsEnergy Levels in Atoms

Energy Levels in Atoms

Objective


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• It explained only a few simple properties of atoms.

Limitations of Rutherford’s Atomic Model


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• It could not explain the chemical properties of elements.


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• It could not explain the chemical properties of elements.

For example, Rutherford’s model could not explain why an object such as the iron scroll shown here first glows dull red, then yellow, and then white when heated to higher and higher temperatures.



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1913, Niels Bohr develops a new atomic model

Bohr stated that the electrons orbit the nucleus like the planets orbit the sun.


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Each possible electron orbit in Bohr’s model has a fixed energy.

•The fixed energies an electron can have are called energy levels.

•Each energy level further from the nucleus is of greater energy


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There are 7 different energy levels

• Each energy level can contain a different amount of electrons


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Niels BohrNiels Bohr’’s s Model (1913)Model (1913)

Electrons orbit the nucleus in circular paths of fixed energy (energy levels).

n=1 first energy level

n=2 second energy lev

n=3 third energy level


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Highest energy level for carbon is n = 2 (2 rings).

Valence electrons – electrons in the outermost energy level


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Bohr's model:-electrons orbit the nucleus like planets orbit the sun-each orbit can hold a specific maximum number of electrons


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The Rutherford model could not explain why elements that have been heated to higher temperatures give off different colors of light.

The Bohr model explains how the energy levels of electrons in an atom change when the atom emits light.


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Ground State: the lowest energy state of an atom.

- An electron absorbs energy (photon) and moves from the ground state to an excited state.

Excited State: when an atom contains excess energy (has higher potential energy).

When an excited atom returns to ground state it gives off light (the energy it has gained as electromagnetic radiation).

Example: Neon signs

Energy and AtomsEnergy and Atoms


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An electron absorbs energy (photon) and moves from the ground state to an excited state.

E4

E3

E2

E1

AbsorptionAbsorption


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When an electron in the excited state returns to the ground state it emits a photon.

E4

E3

E2

E1

Ephoton

= h =E3

-E1

What goes up…must come down!What goes up…must come down!EmissionEmission


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Emitting photons creates light or electromagnetic radiationElectromagnetic radiation in the visible light spectrum has color!These photons have wavelengths that correspond to their color.

Absorption and EmissionAbsorption and Emission


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Unfortunately, Bohr’s model only applied to hydrogen atoms and did not apply to other atoms.

That led scientists to question his model


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Wave Mechanical ModelWave Mechanical Model

Today, the modern description of electrons in atoms is called the Quantum Mechanical Model.

•The wave model tells you the probability of finding an electron in an atom (the exact path of an electron is not known)


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Do Now

What does the Bohr model say about electrons?

Why is this model incorrect?


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There are 7 different energy levels

• Each energy level can contain a different amount of electrons

• There are 4 different types of sublevels


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Energy levels (n=1, n=2 ….)

Sublevels (s,p,d,f)

orbitals


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The sublevels each can contain a different amount of electrons

• s – 2 electrons

• p – 6 electrons

• d – 10 electrons

• f – 14 electrons


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• An atomic orbital is a region of space in which there is a high probability of finding an electron.

• These orbitals have different shapes


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Energy level 1 has only an s sublevel – total

of 2 e-

Energy level 2 has the s and p sublevels – total of 8 e-

Energy level 3 has the s, p, and d sublevels – total of 18 e-

Energy level 4 has the s, p, d, and f sublevels – total of 32 e-


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Do NowDo Now1. How many principal energy levels are there? ____

2. What are the four sublevels? _____

3. How many orbitals does the f sublevel hold? _____

4. How many electrons can each orbital hold? _____


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Do Now AnswersDo Now Answers1. How many principal energy levels are there? 7 (n=7)

2. What are the four sublevels? s, p, d , f

3. How many orbitals does the f sublevel hold? 7

4. How many electrons can each orbital hold? 2


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Aufbau Principle

Electron ConfigurationsElectron Configurations

According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram, each box represents an atomic orbital.

Incr

ea

sin

g e

ne

rgy

6s

5s

4s

3s

2s

1s

6p

5p5d

4p

4d

4f

3p

3d

2p

5.2 Electron arrangement



Hund’s Rule

According to Hund’s rule, electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.

Electron ConfigurationsElectron Configurations5.2 Electron arrangement



Hund’s Rule

Three electrons would occupy three orbitals of equal energy as follows.




Hund’s Rule

Three electrons would occupy three orbitals of equal energy as follows.

Electrons then occupy each orbital so that their spins are paired with the first electron in the orbital.



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Electron configuration describes the placement of the electrons

• Example: Hydrogen: 1s1



Electron ConfigurationsElectron Configurations

Look at the orbital filling diagram of the oxygen atom.

Electron Configurations of Selected Elements

Element 1s 2s 2px 2py 2pz 3sElectron

configuration

H 1s1

He 1s2

Li 1s22s1

C 1s22s22p2

N 1s22s22p3

O 1s22s22p4

F 1s22s22p5

Ne 1s22s22p6

Na 1s22s22p63s1

• An oxygen atom contains eight electrons.










What gives gas-filled lights their colors?

An electric current passing through the gas in each glass tube makes the gas glow with its own characteristic color.

5.3 Atomic Emission Spectra


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By the year 1900, there was enough experimental evidence to convince scientists that light consisted of waves.

The wavelength, represented by (the Greek letter lambda), is the distance between the crests.

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5.3 Electromagnetic Radiation and Energy5.3 Electromagnetic Radiation and Energy


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The frequency, represented by (the Greek letter nu), is the number of wave cycles to pass a given point per unit of time.

The SI unit of waves per second is called the hertz (Hz).



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The frequency () and wavelength () of light are inversely proportional to each other.

As the wavelength increases, the frequency decreases.

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According to the wave model, light consists of electromagnetic waves.

Electromagnetic radiation - a form of energy that exhibits wavelike behavior as it travels through space.

– All electromagnetic radiation travels at the speed of light:

c = 3.0 X108 m/s

Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays.



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Rutherfo

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https://www.youtube.com/watch?v=cfXzwh3KadE


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Equation relating frequency and wavelength:c =

c = speed of light (m/s) = wavelength (m)

= frequency (Hz or s-1)

= c = c

c is constant, so is , so as frequency increases, wavelength decreases (inversely proportional).

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Wave Description of LightWave Description of Light


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c =

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Light as a Wave: ProblemsLight as a Wave: Problems

1) If c = 3.00 x 108 m/s and = 1 x 1019s-1 ,what does equal?

2) What is the frequency of light () if its wavelength () is 4.34 X 10-7 m?


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Visible light of different wavelengths can be separated into a spectrum of colors.

In the visible spectrum, red light has the longest wavelength and the lowest frequency.

Violet light has the shortest wavelength and the highest frequency.


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Atomic Emission Spectrum- a beam of light separated into a series of specific frequencies (and therefore specific wavelengths) of visible light.

– produced when electrons fall back to ground state

– the energy emitted in the fall give off specific patterns (colors) of light

Atomic Emission SpectrumAtomic Emission Spectrum


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The Hydrogen-Atom The Hydrogen-Atom Line Emission Line Emission SpectrumSpectrumThe Hydrogen-Atomic Emission Spectrum


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The Hydrogen-Atomic Emission SpectrumThe Hydrogen-Atomic Emission Spectrum


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Atomic Emission Spectrum of Na, He, Ne, and MercuryAtomic Emission Spectrum of Na, He, Ne, and Mercury


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A fluorescent lamp or a fluorescent tube is a low pressure mercury-vapor gas-discharge lamp that uses fluorescence to produce visible light. An electric current in the gas excites mercury vapor which produces short-wave ultraviolet light that then causes a phosphor coating on the inside of the bulb to glow.



- German physicist Max Planck (1858–1947) showed mathematically that the amount of radiant energy (E) of a single quantum absorbed or emitted by a body is proportional to the frequency of radiation ().

- h is Planck’s constant = 6.63 x 10-34 Js

The Quantization of Energy

E = hEnergy

Planck’s constant (h)= 6.63 x 10-34 Js

Frequency ()


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1. If the frequency () = 1.15 x 1012 s-1 , what is the energy of the radiation?

1. What is the energy of a photon of microwave radiation with a frequency of 3.20 × 1011s-1?

Energy Problems

E = h h = 6.63 x 10-34 Js


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PhotonsPhotons

Einstein proposed that light could be described as quanta of energy that behave as if they were particles.

These light quanta are called photons.


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Key ConceptsKey Concepts

Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus.

The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus of an atom.

Each energy sublevel corresponds to one or more orbitals of different shapes, which describe where the electron is likely to be found.


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Glossary TermsGlossary Terms

• energy level: the specific energies an electron in an atom or other system can have

• quantum: the amount of energy needed to move an electron from one energy level to another


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Glossary TermsGlossary Terms

• quantum mechanical model: the modern description, primarily mathematical, of the behavior of electrons in atoms

• atomic orbital: a mathematical expression describing the probability of finding an electron at various locations; usually represented by the region of space around the nucleus where there is a high probability of finding an electron


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Bohr's Model of the Atom

e.g. fluorine:#P =

#e- =

#N =


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e.g. fluorine:#P = atomic #

= 9#e- =

#N =


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e.g. fluorine:#P = 9

#e- = # P = 9

#N =


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#e- = 9

#N = atomic mass - # P = 10


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#e- = 9

#N = 10

draw the nucleus with protons & neutrons

9P10N


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#e- = 9

#N = 10

how many electrons can fit in the first orbit?

9P10N


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#e- = 9

#N = 10

how many electrons can fit in the first orbit?2

9P10N


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#e- = 9

#N = 10

how many electrons are left?

9P10N


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#e- = 9

#N = 10

how many electrons are left?7

9P10N


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#e- = 9

#N = 10


how many electrons fit in the second orbit?

9P10N


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#e- = 9

#N = 10


how many electrons fit in the second orbit?8

9P10N


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#e- = 9

#N = 10

9P10N


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#e- = 9

#N = 10

How many valence electrons?7

9P10N

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5.1 Revising the Atomic Model > 1 Chapter 5 Electrons In Atoms 5.1 Revising the Atomic Model 5.2 Electron Arrangement in Atoms 5.3 Atomic Emission Spectra