Lab 9

Atomic Structure Emission Spectrum

Electron Configuration

HISTORY OF THE ATOMHISTORY OF THE ATOM

460 BC

Democritus develops the idea of atoms

he pounded up materials in his mortar

and pestle until he had reduced them to

smaller and smaller particles which he

called

ATOMSATOMS

(greek for indivisible)

Development of the Model of an Atom

400 BC Democritus – Particle Model

Ancient Greek philosopher Proposes that matter is composed of smallest

particles called atoms An idea with no evidence

Aristotle – Continuous Model Alternative idea Continuous model of matter- no smallest piece

HISTORY OF THE ATOMHISTORY OF THE ATOM

1808 John Dalton

suggested that all matter was made up of

tiny spheres that were able to bounce

around with perfect elasticity and called

them

ATOMSATOMS

Dalton proposes Atomic Theory in 1803 Based upon experimental evidence All matter composed of atoms Atoms of same element have same mass

and properties Atoms are neither created nor destroyed,

they are rather simply rearranged in chemical reactions Evidence – Lavoiser- Conservation of Mass

Compounds are composed of elements in simple whole number ratios Evidence – Proust

HISTORY OF THE ATOMHISTORY OF THE ATOM

1898 Joseph John Thompson

found that atoms could sometimes eject a

far smaller negative particle which he

called an

ELECTRONELECTRON

Thomson and the Discovery of Electrons

J. J. Thomson’s Experiment

Devised an experiment to find the ratio of the cathode ray particle’s mass (me) to the charge (e)

me /e = –5.686 x 10–12 kg C–1

Thomson’s Plum Pudding Model

Based upon the charge to mass ratio, the electron must be much smaller than the atom

Proposed negative electrons embedded in positive matrix of atom

HISTORY OF THE ATOMHISTORY OF THE ATOM

Thomson develops the idea that an atom was made up of

electrons scattered unevenly within an elastic sphere

surrounded by a soup of positive charge to balance the

electron's charge

1904

PLUM PUDDING

MODEL

Millikan’s Oil Drop Experiment

Measuring the Charge on an electron

Unstable Atoms and Radioactivity

Atoms are not indestructible Atoms are composed of smaller

particles Alpha particles – positively charged Beta particles – negatively charged Gamma rays – no charge

Radioactivity

Goldstein’s Discovery of Protons

Mass Spectrometer- Determining the Percent Abundance of Different Isotopes of same element

Mass Spectrometer

If a stream of positive ions having equal velocities is brought into a magnetic field, the lightest ions are deflected the most, making a tighter circle

Mass Spectrometry

EOS

A record of the separation of ions is called a mass spectrum

Isotopes of Neon Neon-20

10 protons 10 neutrons

Neon-21 10 protons 11 neutrons

Neon-22 10 protons 12 neutrons

HISTORY OF THE ATOMHISTORY OF THE ATOM

1910Ernest Rutherfordoversaw Geiger and Marsden carrying out

his famous experiment.

They fired Helium nuclei at a piece of gold

foil which was only a few atoms thick.

They found that although most of them

passed through. About 1 in 10,000 hit

Rutherford’s Gold Foil Experiment

Rutherford’s Gold Foil ExperimentRutherford’s Gold Foil Experiment

gold foil

helium nuclei

They found that while most of the helium nuclei passed

through the foil, a small number were deflected and, to

their surprise, some helium nuclei bounced straight back.

Rutherford’s Nuclear Model of the Atom

Most the alpha particles (helium nuclei) pass through the gold foil Atom mostly empty space Alpha particles did not hit anything

A very few deflected straight back Alpha particles deflected by a dense

positive nucleus

Visualizing the Pathway of Alpha Particles through a Gold Atom

ELECTROMAGNETIC ELECTROMAGNETIC RADIATIONRADIATION

Electromagnetic radiation.

Electromagnetic RadiationElectromagnetic RadiationElectromagnetic RadiationElectromagnetic Radiation

Most subatomic particles Most subatomic particles behave as PARTICLES and behave as PARTICLES and obey the physics of waves.obey the physics of waves.

Electromagnetic Radiation

The Electromagnetic Spectrum

Wave Model of Light

Wavelength measured in meters (m)

Frequency ƒ measured in waves per second (Hz)

Energy E measured in Joules (J)

Wavelength, Frequency and Energy

Wavelength inversely related to frequency

Increasing wavelength decreasing frequency Decreasing wavelength increasing frequency Wavelength x frequency = speed of light c = ƒ c = speed of light = 3.00 x 108 m/s

Frequency directly related to energy Increasing frequency increasing energy Energy = Planck’s constant x frequency E = h ƒ h = Planck’s constant = 6.63 x 10-34 J/Hz

wavelength Visible light

wavelength

Ultaviolet radiation

Amplitude

Node

Electromagnetic RadiationElectromagnetic RadiationElectromagnetic RadiationElectromagnetic Radiation

ElectroElectromagneticmagnetic SpectrumSpectrumElectroElectromagneticmagnetic SpectrumSpectrum

In increasing energy, RIn increasing energy, ROOYY GG BBIIVV

Sunlight viewed through a spectroscope

Prisms and diffraction grating

bend light

Red light with longer wavelengths bend less

Violet light with shorter wavelengths bend more

Separates white light into ROYGBIV Red – Orange – Yellow – Green – Blue – Indigo – Violet

long waves short waveslow frequency high frequencylow energy high energy

Photoelectric Effect

Photoelectric Effect

Bright red light shined on the photocell has no effect- regardless of intensity or time

Dim green light shined on the photocell causes electrons to be emitted and flow through the wire

Brighter green light emits more electrons per second- greater current

Explaining the Photoelectric Effect ONLY the photon or particle model can

be used to explain these results Energy is required to pull off negative

electrons attracted to positive protons in nucleus- breaking attraction

Each red photon does not have enough energy to pull off an electron – regardless of how long the light is shined

Each green photon has more energy and can pull off the electron when the green photon collides

Light Spectrum Lab!

Slit that Slit that allows light allows light insideinside

Line up the slit so Line up the slit so that it is parallel with that it is parallel with the spectrum tube the spectrum tube (light bulb)(light bulb)

The Emission Spectrum of Hydrogen- Discrete Bands of Colored Light

Excited Gases Excited Gases & Atomic & Atomic StructureStructure

Emission Spectra of Different Atoms: A Fingerprint to Identify

Rydberg and Balmer (1886)

Independently develop mathematical equations that fit the data for hydrogen emission spectrum

The electron had no yet been discovered

Neither had a model to explain the observed wavelengths

Just an equation that worked

Rydberg Equation

1/λ = RH [1/n12 - 1/n2

2]

RH = 1.09678 x 10-2 nm-1

Solve the equation for an electron

moving from level 4 to 2.

Bohr’s Model of the Atom (1910)

Assumed electrons orbit the nucleus in circular orbits

Proposed the energy of the orbit is proportional to the distance from the nucleus (increasing distance – increasing energy)

Assumed only certain allowable energies

Used angular momentum to calculate the allowable energy

Bohr’s Model of the Atom (1910)

When the atom absorbs energy Electron moves up to higher energy with

more potential energy farther away from nucleus

Unstable with higher PE Electron falls back down to lower levels

Energy released PE converted to KE as electron fall The color of light observed reflects the

energy released in the fall

Bohr’s Calculations

of the Energy

ΔE = -2.18 x 10-18 J (1/nf2 – 1/ni

2)

n = the energy level

ΔE = positive when electron climbs up levels

absorbing energy

increasing PE

ΔE = negative when e- falls down levels

releasing energy

decreasing PE

Niels BohrNiels Bohr

(1885-1962)(1885-1962)

ΔE = -2.18 x 10-18 J (1/nf2 – 1/ni

2)

Calculate the energy as an electron drops from level 6 down to level 2.

Calculate the frequency and wavelength of this photon.

ultraviolet infrared

Visualizing the Movement of the

Electron

Line Spectra of Line Spectra of Other ElementsOther Elements

Oops. Bohr’s equation does NOT predict these wavelengths.

Visualizing the “falling” e-

Where does the electron have more potential energy?

Electron is a wave - De Broglie

De Broglie Since light is both particle and wave,

perhaps so is matter both particle and wave Wavelength depends upon mass and

velocity Objects with large mass have negligible, so small we can ignore Electrons with very small mass have wave

properties that cannot be ignored

λ = h/p

Quantum Mechanics Heisenberg

Uncertainty principle Given the wavelike nature of electron Impossible to know both location and energy of

electron Can only calculate the probable location

Schrodinger Used calculus to “locate” the electron within

orbital Orbitals – regions of space representing the most

probable location of electron

E. SchrodingerE. Schrodinger1887-19611887-1961

W. HeisenbergW. Heisenberg1901-19761901-1976

Wave Functions: Calculating the Probability of locating an electron in a region of space

The Uncertainty Principle: Cannot both determine location and energy of electron

The Wave Function and Orbitals

The region near the nucleus is separated from the outer region by a spherical node - a spherical shell in which the electron probability is zero

EOS

Quantum Numbers

Values that emerge from the wave functions of Schrodinger

1st n = energy level 2nd ℓ = shape of orbital (s, p, d or f) 3rd mℓ = orientation (diff

versions) 4th ms

= magnetic spin of electron

Increasing Radius of s-orbital with higher values of n

s orbitals orbital p orbitalp orbital d orbitald orbital

f Orbitalsf Orbitalsf Orbitalsf Orbitals

The s-orbital

Spherical shaped orbital

ℓ = 0

mℓ = 0

Only one s-orbital in any energy level

The p-orbitalDouble-lobe shaped orbitalℓ = 1mℓ = -1 or 0 or +1Only three p-orbitals in any energy level- except for level one

Planes of zero probability

Model of d-orbital

Only electrons with opposite spins can be in the same orbital

Electron Configurations

Show the electrons in orbitals Box used to represent orbital Half arrow used to represent e- with opposite spins

Electrons are placed in orbitals of lowest energy first

Use sum of first two quantum numbers to determine which orbital fills first

1s

2s

3s3p

2p

1s

2s

3s3p

2p

Nickel Electron Configuration and quantum numbers

1s 2s 2p 3s 3p

1 0 0 ½ 2 0 0 ½ 2 1 -1 ½ 2 1 0 ½ 3 0 0 ½ 3 1 0 ½

2 1 0 ½ 3 1 -1 ½ 3 1 0 ½

4s 3d

4 01 0 ½ 3 2 -2 ½ 3 2 -1 ½ 3 2 0 ½

3 2 1 ½ 3 2 1 ½

1st # indicates energy level

n = 1 1st level

n = 2 2nd level

2nd # type of orbital

l = 0 is s-orbital

l = 1 is p-orbital

l = 2 is d-orbital