7 - 1 Atomic Structure The Electron The Nucleus Traveling Waves Electromagnetic Radiation Bohr Model...

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Atomic StructureAtomic Structure

The ElectronThe ElectronThe NucleusThe Nucleus

Traveling WavesTraveling WavesElectromagnetic Radiation Electromagnetic Radiation

Bohr Model of the Hydrogen AtomBohr Model of the Hydrogen AtomWave Theory of the ElectronWave Theory of the Electron

Heisenberg Uncertainty PrincipleHeisenberg Uncertainty PrincipleQuantum Model of the AtomQuantum Model of the Atom

Discovery of the electronDiscovery of the electron

18071807 Davy suggested that electrical forces held compounds together.

18331833 Faraday related atomic mass and the electricity needed to free an element during electrolysis experiments.

18911891 Stoney proposed that electricity exists in units he called electrons.

18971897 Thomson first quantitatively measured the properties of electrons.

Cathode raysCathode rays

Thomson’s ‘discovery’ of electrons was based on studies of cathode rays.

These are produced when a gas is ionized.

Ionized gas

cathode anode Smalldot

Thomson observed that the position of the dotwas altered when either electrical or magneticfields were applied.

Thomson was not able to measure the mass or charge of an electron.

He could only determine the mass to charge ratio -- 6 x 10-12 kg/C

Columb (C) - the SI unit of charge.

It is defined as the amount of charge that flows past a fixed point in a wire per second when the current is one ampere.

Charge of the electronCharge of the electron

Millikan experimented with electrically charged oil drips.

By varying the electrical field, he observed that the drops had charges that were whole number multiples of 1.5924 x 10-19C

This represented the charge of an electron.

The modern value is now known to be -1.602 1773 x 10-19C

Mass of the electronMass of the electron

Once both the mass to charge ratio and the actual charge of an electron were known, finding the mass was pretty easy.

me = ( mass to charge ratio ) ( charge )

= ( 6 x 10-12 kg C-1 ) ( 1.5924 x 10-19 C)

= 1 x 10-30 kg

The modern value for the mass of an electron is: 9.109 390 x 10-31 kg

5.485 799 x 10-4 u

Discovery of the nucleusDiscovery of the nucleus

19091909 Rutherford bombarded thin metal foils with alpha particles (helium ions).

He felt that the particles would pass through the foils.

When tested, he observed that about 1 alpha particle in 8000 was deflected by the foil.

This deflection indicated the existence of a small, dense, positively charged nucleus.

Discovery of the nucleusDiscovery of the nucleus

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Determination of nuclear chargeDetermination of nuclear charge

Rutherford estimated that the charge of the nucleus of an atom was about one half of the atomic mass.

Moseley, while working for Rutherford, developed a more accurate measurement.

While working with cathode rays, he measured the wavelength of the X-rays produced.

He found that a direct relationship exists between the atomic number and the square root of the frequency.

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Determination of nuclear chargeDetermination of nuclear charge

Moseley concluded thatthe charge of the nucleuswas an integer.

Further, it was the sameas the number of electricalunits (electrons) but ofopposite charge.

Moseley concluded thatthe charge of the nucleuswas an integer.

Further, it was the sameas the number of electricalunits (electrons) but ofopposite charge.

X-Ray Frequency1/2

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Discovery of protons and neutronsDiscovery of protons and neutrons

Measurements of the mass to charge ratio of the nucleus were made in a manner similar to earlier work with electrons.

The ratio was found to be dependent on the gas that was used for the experiment.

Hydrogen was found to produce particles with the lowest mass. These particles were assumed to be common to all atoms and were called protons.

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Discovery of protons and neutronsDiscovery of protons and neutrons

19321932 Chadwick observed that when beryllium-9 was exposed to alpha particles, particles with the same mass as protons but no charge were given off.

These were called neutrons and are present in all atoms except

hydrogen-1.

They contribute to the force that holds the nucleus together and reduce the

repulsive force between positively charged protons.

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Traveling wavesTraveling waves

Much of what has been learned about atomic structure has come from observing the interaction of visible light and matter.

An understanding of waves and electromagnetic radiation would be helpful at this point.

Let’s start with some basic definitions.

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WavesWaves

Some definitions

Wavelength, Wavelength, The distance for a wave to go through a complete cycle.

AmplitudeAmplitudeHalf of the vertical distance from the top to the bottom of a wave.

Frequency, Frequency, The number of cycles that pass a point each second.

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WavesWaves

wavelength

+ +- - -

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Electromagnetic radiationElectromagnetic radiation

A form of energy that consists of perpendicular electrical and magnetic fields that change, at the same time and in phase, with time.

The SI unit of frequency () is the hertz, Hz

1 Hz = 1 s1 Hz = 1 s-1-1

Wavelength and frequency are related

= = cc

c is the speed of light, 2.998 x108 m/s

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1020 1015 1010 105 100

10-10 10-5 100 105 1010

Wavelength (), m

Frequency (), s-1

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Separation of lightSeparation of light

‘White’ light is actually a blend of all visible wavelengths. They can separated using a prism.

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Electromagnetic radiation (EM) and matterElectromagnetic radiation (EM) and matter

TransmissionTransmission - EM will pass through matter -- no interaction.

AbsorptionAbsorption - EM is absorbed by an atom, ion or molecule, taking it to a higher energy state.

EmissionEmission - the release of energy by an atom, ion or molecule as light, taking it to a lower energy state.

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Particle propertiesParticle properties

Although EM has definite wave properties, it also exhibits particle properties.

Photoelectric effect.Photoelectric effect.•First observed by Hertz and then later

explained by Einstein.•When light falls on Group IA metals,

electrons are emitted (photoelectrons).•As the light gets brighter, more

electrons are emitted. •The energy of the emitted electrons

depends on the frequency of the light.

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Photoelectric effectPhotoelectric effect

The cathode has a photoemissive surface.

When light hits the cathodeelectrons are ejected.

They are collected at the anode and can bemeasured.

cathode

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Photoelectric effectPhotoelectric effect

Studies of this effect led to the discovery that light existed as small particles of electromagnetic radiation called photonsphotons.

The energy of a photon is proportional to the frequency.

Photon energy = Photon energy = hhThe energy is inversely proportional to the wavelength.

Photon energy = Photon energy = hh c c -1-1

h - Plank’s constant, 6.626 x 10-34 J . s

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Photon energy examplePhoton energy example

Determine the energy, in kJ/mol ofa photon of blue-green light with a wavelength of 486 nm.

energy of a photon =

= 4.09 x 10-19 J / photon

(6.626 x 10-34 J.s)(2.998 x 108 m.s-1)(4.86 x 10-7 m)

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Photon energy examplePhoton energy example

We now need to determine the energy for a mole of photons (6.022 x 1023)

Energy for a mole of photons.Energy for a mole of photons.

= (4.09 x 10-19 J / photon) (6.022 x 1023

photons/mol)

= 246 000 J/mol

Finally, convert to kJ

= ( 244 000 J/mol )

= 244 kJ / mol

1 kJ103 J

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Bohr model of the atomBohr model of the atom

Bohr studied the the spectra produced when atoms were excited in a gas discharge tube.

He observed that each element produced itsown set of characteristic lines.

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Balmer later determined an empirical relationship that described the spectral lines for hydrogen.

1 = 1.097 x 107 m-1( )1

n = 2, 3, 5, . . .

Spectra of many other atoms can be described bysimilar relationships.

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Bohr proposed a model of how electrons moved around the nucleus.

• He wanted to explain why electrons did not fall in to the nucleus.

• He also wanted to account for spectral lines being observed.

He proposed that the energy of the electron was quantized - only occurred as specific energy levels.

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In the Bohr model, electrons can only exist at specific energy levels (orbit).

Each energy level was assigned a principal quantum number, n.

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The Bohr model is a ‘planetary’ type model.

Each principal quantum represents a new ‘orbit’ or layer.

The nucleus is at the center of the model.

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Bohr was able to use his model hydrogen to:Bohr was able to use his model hydrogen to:

• Account for the observed spectral lines.

• Calculate the radius for hydrogen atoms.

His model did not account for:His model did not account for:

• Atoms other than hydrogen.

• Why energy was quantized.

His concept of electrons moving in fixed orbits was later abandoned.

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Wave theory of the electronWave theory of the electron

19241924 De Broglie suggested that electrons have wave properties to account

for why their energy was quantized.

He reasoned that the electron in the hydrogen atom was fixed in the

space around the nucleus.

He felt that the electron would best be represented as a standing wave.

As a standing wave, each electron’s path must equal a whole

number times the wavelength.

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De Broglie proposed that all particles have a wavelength as related by:

= wavelength, metersh = Plank’s constantm = mass, kgv = frequency, m/s

De Broglie wavesDe Broglie waves

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De Broglie wavesDe Broglie waves

Using De Broglie’s equation, we can calculate the wavelength of an electron.

=6.6 x 10-34 kg m2 s-1

(9.1 x 10-31 kg)(2.2 x 106 m s-1)

The speed of an electron had already been reportedby Bohr as 2.2 x 106 m s-1.

= 3.3 x 10-10 m

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Heisenberg uncertainty principleHeisenberg uncertainty principle

• In order to observe an electron, one would need to hit it with photons having a very short wavelength.

• Short wavelength photons would have a high frequency and a great deal of energy.

• If one were to hit an electron, it would cause the motion and the speed of the electron to change.

• Lower energy photons would have a smaller effect but would not give precise information.

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Heisenberg uncertainty principleHeisenberg uncertainty principle

According to Heisenberg, it is impossible to know both the position and the speed of an object precisely.

He developed the following relationship:

As the mass of an object gets smaller, the product of the uncertainty of its position and speed increase.

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Quantum model of the atomQuantum model of the atom

Schrödinger developed an equation to describe the behavior and energies of electrons in atoms.

• His equation is similar to one used to describe electromagnetic waves.

• While the equation is too complicated to write here, we can still use the results.

• Each electron can be described in terms of its quantum numbers.

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Quantum numbersQuantum numbers

Principal quantum number, Principal quantum number, n nTells the size of an orbital and largely determines its energy.

n = 1, 2, 3, ……

Angular momentum, Angular momentum, llThe number of subshells that a principal level contains. It tells the shape of the orbitals.

l = 0 to n - 1

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Magnetic quantum number, mMagnetic quantum number, mll

Describes the direction that the orbital projects in space.

ml = l to +l (all integers, including zero)

For example, if l = 2, then ml would have values of -2, -1, 0, 1 and 2.

Knowing all three numbers provide us with a picture of all of the orbitals.

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subshell # ofn l ml label orbitals

1 0 0 1s 1

2 0 0 2s 11 -1, 0, 1 2p 3

3 0 0 3s 11 -1, 0, 1 3p 32 -2, -1, 0, 1, 2 3d 5

4 0 0 4s 11 -1, 0, 1 4p 32 -2, -1, 0, 1, 2 4d 53 -3, -2, -1, 0, 1, 2, 3 4f 7

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The The ss orbital orbital

The s orbital is a sphere. Every level has one s orbital.

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pp orbitals orbitals

There are three p orbitals: px, py and pz

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Representative Representative dd orbitals orbitals

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Representative Representative ff orbitals orbitals

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Combined orbitals - Combined orbitals - nn=2=2

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Combined orbitals - Combined orbitals - nn=3=3

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Electron spinElectron spin

Pauli added one additional quantum number that would allow only two electrons to be in an orbital.

Spin quantum number, Spin quantum number, mmss..It can have values of +1/2 and -1/2

Pauli also proposed that no two electrons in an atom can have the same set of four quantum numbers --

Pauli exclusion principlePauli exclusion principle.

7 - 1 Atomic Structure The Electron The Nucleus Traveling Waves Electromagnetic Radiation Bohr Model...

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