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Chapter SevenAtomic Structure

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The Electron: History of Cathode Rays1838: Michael Faraday

Electricity flows causes purple glow in a tube at near vacuum Few molecules: minimal molecular interferenceGlow from absorption of current by remaining air molecules

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Numerical Values for Atomic Particles

1897: J.J. ThomsonRatio of cathode ray particle’s mass to charge

me / e- = –5.686 X 10-12 kg\CNegatively charged particle is called an electron

Robert MillikanValue for the charge on an electron

e = -1.602 X 10-19 C

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Millikan’s Oil Drop ExperimentAtomize oil

Tiny, negative oil dropletsMeasure of velocity under gravityTurn on electric field

Velocity decreasesAttraction to (+) pole

Calculate particle chargeMultiples of same numbere- = - 1.602 X 10-19 C

Calculate electron mass Charge to mass ratio me / e- = –5.686 X 10-12 kg\CElectron charge e- = - 1.602 X 10-19 CElectron Mass emass = 9.109 X 10-31 kg\ electron

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Positively charged sphere with electrons imbedded inside“Raisin pudding model”

Hydrogen atom1 electron at the center of a sphere

Helium atom:2 electrons along a straight line

Applied this analysis to atoms with up to 100 electrons

The Atom: J. J. Thomson’s Model

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Rutherford’s ExperimentsTested JJ Thomson’s theory of electrons in center of atom

Bounced alpha particles (He2+) through gold foilMost went through, some slightly deflected,

Some bounced back completely

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Rutherford’s ConclusionsDetermined Thomson’s theory flawed and concluded new model

NucleusHeavyDense(+) charged (He2+)

Electron “Cloud”Light and fluffySurrounds nucleus(- ) charged

Experimental Resultsα particles that go all the way through hit nothingα particles deflected hit electrons in electron cloudα particles bouncing back hit nucleus

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Final ModelPositive charge carried by particles called protons

Proton charge: fundamental unit of positive chargeThe nucleus of a hydrogen atom consisted of 1 proton

Atomic number# of protons in the nucleus

NeutronsDiscovered by James ChadwickAtomic particle with nearly the same mass as protonsNo charge- accounts for extra mass of isotopesChange in mass, no change in properties

+ -Hydrogen(Deuterium)

electron

proton neutron

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Electromagnetic RadiationWave

Repeating disturbance spreading out from a defined origin

Electromagnetic waves Originate from charge movementProduce fluctuations in electric and magnetic fieldsRequire no medium

Characterized by wavelength, frequency, and amplitudeWavelength: Distance between 2 identical points in a cycleFrequency: The number of times the repetition passes a

given point per secondAmplitude: Height of wave from center to peak

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WavelengthDistance between 2 identical points in a cycle

Symbolized by λ

Units of length depend on range used Nanometers (nm) = 10-9 m

Visible light: 390 nm to about 760 nm.Angstrom (A) = 10–10 m

Visible light is 3900 A to 7600 A.

λ

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FrequencyThe number of times the repetition passes a given point

per second

Symbolized by νHertz (Hz) = 1 cycle per second = 1 cycle/sMeasurement of pitch of a sound wave

λ

λ

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Relationship of Wavelength and Frequency

Wavelength x Frequency = Constant = Speed of light = c

c =λ x ν =3.0 x108 m/sLong wavelength: Lower frequencyShort wavelength: Higher frequency

λ

λ

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What is the wavelength of an FM-radiowavewith a 94.9 MHz frequency?

c = λν =3.00 x 108 m/sλ = c/ν

94.9M Hz = 94.9 x 106 Hz = 94.9 x 106/s

mxsx

smx 16.3

109.9411000.3

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=⎥⎦⎤

⎢⎣⎡

⎥⎦

⎤⎢⎣

⎡=λ

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Electromagnetic SpectrumComplete range of wavelengths and frequenciesExtends continuously from shortest wavelengths to longest.

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Continuous and Line SpectraContinuous spectrum:

Wavelengths mergeElectromagnetic spectrum is continuousWhite light: Consists of all colors of the spectrum.

Line spectrum:Occurs when light is produced through an elementSpectrum is discontinuous

A pattern of lines characteristic of that elementUsed for Identification

Emission Spectroscopy is the analysis of light emitted from a strongly heated or energized element

Emission Spectrum: A record of the emitted light16

Emission Spectrum of Helium

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Planck’s ConstantMax Planck’s Hypothesis

Energy is absorbed or emitted in discrete amounts.May be in multiples of these discrete amounts

Quantum: The smallest amount of energy: E = hνPlanck’s constant: h = 6.626 X 10-34 J s

EinsteinEnergy bundled into little packetsEnergy of 1 packet= 1 photon= 1 quantum of energyEphoton = hν

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The Photoelectric EffectA full spectrum beam of light hits a metal surface Energy transferred to electrons in metal creating currentElectrons in metal break attraction to metal ions and escape The energy leaves in discrete packets, not as a single beam Einstein called these packets “photons”.

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Energy Calculations

Einstein: Energy is determined by its frequencyEnergy of photon = E = hν

Energy/mol = NA (photons/mol) X h ( J s/photon) X ν (s-1)

Use to compute energy of various wavelengths of lightc = λν and E = hν

Ephoton = hc/λ

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ExamplesWhat is the energy of a radiowave with a frequency of 94.9 MHz?

What is the energy per photon and per mole of photons of violet light, with a wavelength of 415 nm?

What wavelength has an energy of E = 1.00 x 10-20J?

JxsxxJsxhEphoton

26634

1023.6100.94110626.6 −

−

=== υ

Jxmx

xs

mxxJsxhcE photon19

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834

1079.41041511000.3

110626.6 −

−

−

===λ

mxJx

xs

mxxJsx 520

834

1099.11000.111000.3

110626.6 −

−

−

==λ

molJx

molphotonsxx

photonJxEmol

52319 1088.21

1002.61079.4=

−

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Energy Level CalculationsAll calculations done by comparing energy levels

Initial level: Ei = - B / ni2

Final level: Ef = - B / nf2

Elevel = -B/nf2 - -B /nf

2 = B (1/ni2 – 1/nf

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Energy is emitted or absorbed when an electron moves between energy levels.High to low level

energy released: atom more stableLow to high level

energy absorbed: atom less stable= 0 at a level infinitely far from the nucleus

Ground state: The lowest possible energy level Excited state: All other levels

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What is the energy of an electron in a n=4 energy level?

JxJxnBE 19

2

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2 10362.14

10179.2 −−

−=−

=−

=

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What is the wavelength of the shift from n = 4 to n =2? Is light emitted or absorbed?

⎟⎟⎠

⎞⎜⎜⎝

⎛−=⎟⎟

⎠

⎞⎜⎜⎝

⎛ −−

−= 2222

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fiif nnB

nB

nBE

JxJxxE 1922

18 10086.421

4110179.2 −− −=⎟

⎠⎞

⎜⎝⎛ −=

so E Ehchc

== λλ

mxJx

xs

mxxJx 719

834

1086.410086.4

11000.310626.6 −−

−

==λ

λ = 486.2 nm Visible blue green light is emitted (higher to lower n)24

Emission SpectraTogether, all the photons making the same change from one

energy level to another produce one spectral line

The collection of lines is the emission spectrum

Emission Spectrum for Mercury, Hg

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Wave Mechanics: Matter as Waves

Light has both wave and particle propertiesWave- light dispersed into a continuous spectrumParticle- Discrete photons displace electrons

Louis de Broglie (1923)Matter can behave as both particles & wavesA particle with a mass, m, moving at a speed, v, will have a

wavelength (λ) = h/mvPrediction of matter waves led to the development of the

electron microscope

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Bohr’s Hydrogen AtomNiels Bohr (1913): Electron energy (En) was quantized

Only certain specified values allowed (stable orbits)A specified energy value is an energy level of the atomThe energy of each stable orbit En = –B/n2

n is the quantum number for the atom integers only, 1,2 3..B = 2.179 X 10-18 J

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Quantum Numbers and Atomic OrbitalsAtomic orbital

A region in space with a high probability of finding an electron.Identified by 4 quantum numbers.

Four Quantum Numbers1. The principal quantum number (n)2. The orbital quantum number (l)3. The magnetic quantum number (ml)4. The spin quantum number (ms)

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The Principal Quantum Number (n)Restricted to the positive integers: 1, 2, 3, 4, 5,…Indicates the shell or level of the orbitalIndicates the size of the orbitalIntegers correspond to row numbers in periodic table

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2

3

4

5

6

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The Orbital Quantum Number (l)Non-negative integers smaller than n

Designates the subshellIndicates the shape of the orbital

I = 0 are called s orbitalsSpherical

l = 1 are called p orbitals2 teardrops joined point- to- pointReferred to as Px,Py,Pz

I= 2 are d orbitalsl= 3 are f orbitals.

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The Magnetic Quantum Number (ml):Determines the orientation in space of the orbitals

Any integer from - lto +lThe number of possible values for ml = 2l + 1Determines the number of orbitals in a subshell

Corresponds to Px,Py,Pz orbitals of the P subshell

l= 1 # values = 3 Range: -1 to 1ml = -1,0,1

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Electron Spin Quantum Number (ms)Explains some of the finer features of atomic emission spectraThe number can have two values: +1/2 and –1/2The spin refers to a magnetic field induced by the moving

electric charge of the electron as it spinsThe magnetic fields of two electrons with opposite spins cancel

one another; there is no net magnetic field for the pair.

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Quantum Numbers Summary