1. Atomic Structure
Chapter 6

2. 6.1 Electromagnetic Radiation
All forms of radiation (such as light, microwaves, radio) are forms of energy that can be described in a mathematical theory called electromagnetic radiation.
The distance between successive crests of a wave (or between successive troughs) is the wavelength of a wave. Wavelength is symbolized by λ (lambda) and usually measured in meters or nanometers.

3. Wave Properties
Frequency is the number of wavelengths (or cycles) that pass a given point in a second. It is symbolized by ν (nu).
number of events per time
unit of s-1 (1 per second) or Hz (hertz)

4. Wave Properties
Wave height is called amplitude and points of zero amplitude are called nodes. Nodes occur at intervals of λ/2.
Speed of a wave can be calculated by multiplying wavelength times frequency.
C = λν
where C is the speed of light (3.0 x 108 m/s)

5. Electromagnetic Spectrum
The electromagnetic spectrum includes all the wavelengths of radiant energy from short gamma rays to long radio waves.
The visible spectrum is the part of the spectrum visible to the eye, usually with wavelengths between 400 and 700 nm.

6. Example Problem
The wavelength of the radiation which produces the yellow color of a sodium vapor light is nm. What is the frequency of radiation?
Orange light has a _____ frequency and a _____ wavelength than blue light. (see page 271)

7. 6.2 Planck, Einstein, Energy and Photons
Before Planck, predictions to describe the relationship between wavelength and radiation failed in the ultraviolet region.
Max Planck introduced the concept of quantized vibrations. Quantization means that only certain vibrations (specific frequencies) are allowed.

8. Planck, Einstein, Energy and Photons
A quantum (or photon) is a particle of light energy that can be absorbed (or emitted) by a molecule. The energy of that molecule is increased (decreased) by an amount equal to the energy of the photon.
Planck’s constant is a proportionality constant that describes the energy of a photon.
E = hv
where E is energy (J), h is Planck’s constant (J∙s), and v is frequency (s-1)

9. Planck, Einstein, Energy and Photons
Light can cause chemical reactions to occur! (paint fades, cloth decomposes)
As frequency increases, energy of radiation increases. Energy also increases as the wavelength decreases.
E = hv = (hc)/λ
(example: UV light has more energy than visible light due to shorter wavelengths)

10. Planck, Einstein, Energy and Photons
higher intensity of light would mean there are more photons to strike a surface per unit of time
photons have to have enough energy to remove electrons from an atom
once the minimum energy (light frequency) is exceeded, the photons have enough energy to displace electrons
more high energy photons means more electrons displaced

11. ν = c/λ E = hν E = hc/λ λ = wavelength in nm
V = frequency in 1/s or Hz
E = energy of a single photon in joules
C = speed of light = 3.00 x 1017 nm/s
h = Planck’s constant = 6.63 x J-s

12. Planck, Einstein, Energy and Photons
The photoelectric effect occurs when light strikes the surface of a metal and electrons are ejected.
Einstein combined Planck’s equation with the concept of photons, the “particles” of light. Electromagnetic radiation is thought of a stream of photons.
Matter is allowed to emit or absorb energy only in discrete amounts.

13. Practice Problem
What is the frequency, the energy of a single photon, and the energy of a mole of photons of light having a wavelength of 555 nm?

14. Practice Problem
Calculate (a) the energy in joules of a photon emitted by an excited sodium atom (wavelength 600 nm) and (b) the energy in kJ of a mole of photons.

15. Homework
After reading sections 6.1 and 6.2, you should be able to do the following…
p. 297 (3-12)

16. 6.3 Atomic Line Spectra and Niels Bohr
A spectrum that consists of light of all wavelengths is called a continuous spectrum.
A line emission spectrum (or atomic emission spectrum) occurs when light from excited electrons emit only certain wavelengths of light.
Each element produces a characteristic and identifiable pattern.

17. Line Emission Spectra
Hydrogen’s series of 4 lines is referred to as the Balmer series.

18. Bohr Model
Niels Bohr connected the emission spectra with Planck and Einstein’s ideas.
Bohr proposed that electrons move in circular, fixed energy orbits around the nucleus.
Each circular orbit corresponds to a stable energy state.
Bohr introduced quantization into electronic structure!

19. Bohr Model
An atom with electrons in the lowest possible energy levels is said to be in its ground state. When the electron of a hydrogen atom occupies an orbit with n greater than 1, it is said to be excited state.

20. Bohr Model
Energy (a photon) is emitted or absorbed by an electron when it changes from one allowed energy state to another.
The lines of the atomic emission spectrum of hydrogen result when an electron falls from a higher allowed state to a lower allowed state.
The increment between each allowed state is proportional to Planck’s constant, the speed of light, and the Rydberg constant, RH.

21. Atomic Line Spectra
The Rydberg equation allows us to calculate the wavelengths of different lines in the visible emission spectra of hydrogen atoms.
You can use the Rydberg constant and Planck’s constant to calculate the energy states at different levels.
where R = x 107 m-1

22. Practice Problem
Calculate the energies of the n = 3 states of the hydrogen atom in joules per atom and in kilojoules per mole.

23. Bohr Theory and Spectra
Electrons stay in lower energy levels unless they absorb or evolve energy due to some disturbance.
Electrons in the ground state have energy with a large negative value. As the electron absorbs energy and moves to a higher energy level, its energy becomes less negative.
positive change in energy – absorption
negative change in energy – emission

24. Bohr Theory and Spectra
As electrons naturally move back to lower levels, they emit energy which is observed as light.
This is the atomic emission spectrum! The movement of electrons between quantized energy states.
Bohr’s model only applies to H atoms or systems with one electron.

25. 6.4 Particle Wave Duality
Louis Victor de Broglie proposed that matter, such as an electron, could exhibit wavelike properties. He said that an electron with mass (m) moving with velocity (v) should have a wavelength…
λ = h/(mv)
Electrons were experimentally shown to be diffracted (like light waves) by a thin sheet of foil, so therefore electrons can have wave properties.

26. Practice Problem
Calculate the wavelength associated with an electron of mass m = x g traveling at 60.0% of the velocity of light.

27. Homework
After reading sections 6.3 and 6.4, you should be able to do the following…
p. 298 (14-20,23-26)

28. 6.5 Quantum Mechanical View
The general approach to understanding atomic behavior that includes theories of Bohr, Schrodinger, and others is called quantum mechanics or wave mechanics.
Like light, electrons have properties of both a wave and a particle.
Werner Heisenberg concluded that it is impossible to fix both electron position and energy if the electron is described as a wave.

29. Heisenberg Uncertainty Principle
The uncertainty principle, applied to electrons in an atom, states that it is inherently impossible to simultaneously determine exact position and momentum of an electron. The best that can be done is to predict probability of finding an electron in a certain region of space.

30. Schrodinger’s Model
Erwin Schrodinger developed mathematical equations with solutions called wave functions (ψ – psi) that are chemically important.
The square of a wave function is an orbital.
orbital - probability of finding an electron of a given energy in a region of space (electron density)

31. Schrodinger’s Model
The region of space in which an electron of a given energy is most probably located is called its orbital.
Three integer numbers – the quantum numbers n, l, and ml – are an integral part of the mathematical solution.

32. Practice Questions
Who proposed the wave-particle properties of electrons?
Who discovered the charge-mass ratio of an electron?
Who provided the theoretical explanation of the photoelectric effect?
Who first postulated that the sharp lines in the emission spectra of elements were caused by electrons going from high energy levels to low energy levels?

33. Quantum Mechanical Model
The quantum mechanical model of the atom is a mathematical model that incorporates both wave and particle characteristics of electrons in atoms.

34. Homework
After reading section 6.5, you should be able to do the following…
p. 299 (34-38) – needs to be adjusted for quantum numbers!

35. 6.6 Shapes of Atomic Orbitals
Electrons with the highest value of n are valence electrons.

36. s - orbital
electrons cluster around the nucleus most of the time (probabilty is that electrons will be found within that radius about 90% of the time) – electron cloud picture
electron density is greater closer to nucleus
square of wave function (Ψ2) is probability density – high for points around nucleus
s orbital is spherical in shape
the size of s orbitals and their energy increases as n increases

37. p - orbitals
All p orbitals have a nodal surface that slices through the nucleus and divides the region of electron density in half
dumbbell shaped
3 possible orientations

38. d - orbitals
Orientation possibilities are equal to the number of nodal surfaces that slice through the nucleus
2 nodal surfaces divides into four regions of electron density
5 d orbitals

39. f - orbitals
seven orbitals
7 regions of electron density

40. 6.7 Electron Spin
Electrons have an intrinsic property known as spin that can result in atoms having a magnetic moment.
At most, two electrons can be accommodated in an orbital, and these electrons must have opposite spin.

41. Magnetism
Electrons act as “micromagnets”, and there are only two spins possible…
A material that is slightly repelled by a strong magnet is said to be diamagnetic, materials that are attracted to a strong magnet are paramagnetic; they lose their magnetism once removed from the field. Ferromagnetic materials retain magnetism upon introduction to removal from a magnetic field.

42. Magnetism
How can we account for the fact that the H atom is paramagnetic and the He atom is diamagnetic?

43. Atomic Orbitals and Chemistry
By thinking about orbitals of atoms in molecules, and by making simple assumptions that they resemble those of the hydrogen atom, we can understand much of the chemistry of complex systems.