1. Test 5: Chapter 5 – Electrons in Atoms
Chemistry 2009

2. Explore Activity – Mystery Envelopes
Work in small groups of 3-4
Using the procedures for the Discovery Lab on p. 117, determine what is inside your envelope.
Copy your procedures, observations, and analysis.
Turn it in to the bin!

3. Light and Quantized Energy
Section 1:
Light and Quantized Energy

4. The Nuclear Atom & Unanswered Questions
Rutherford’s nuclear model proposed that all of an atom’s positive charge and virtually all of its mass are concentrated in a nucleus that is surrounded by fast-moving electrons. (nuclear model)
Major scientific development
Lacked details about how the electrons occupy the space surrounding the nucleus
Physicists said that it did not explain electron arrangement or how the electrons were not pulled into the nucleus
Chemists said that it did not account for differences in chemical behavior between elements

5. Early 1900’s scientists studied chemical behavior:
Observed that certain elements emitted visible light when heated in a flame
Analysis of the light revealed that behavior is connected to electron structure.

6. Wave Nature of Light
Electromagnetic radiation – a form of energy that exhibits wavelike behavior as it travels through space.
Ex. Visible light, microwaves, x-rays, radio and television waves.
Primary characteristics of waves:
Wavelength
Frequency
Amplitude
Speed

7. Wavelength
Wavelength – the shortest distance between equivalent points on a continuous wave
Represented by λ (Greek letter lamda)
Generally measured crest to crest or trough to trough
Usually expressed in meters, centimeters, or nanometers

8. Frequency
Frequency – the number of waves that pass a given point per second
Represented by ν (Greek letter nu)
SI unit = hertz (Hz) (1Hz = 1 wave per second)
In calculations – frequency is expressed with units of “waves per second”, #waves/second, or s-1
Ex. 652 Hz = 652 waves/second = 652/s = 652 s-1

9. Amplitude
Amplitude – wave’s height from origin to crest or origin to trough

10. Speed
Speed – all waves travel at a speed of 3.00 x 108 m/s in a vacuum
Speed of light symbol = c
c=wavelength x frequency or c = λν
Although the speed of all electromagnetic waves is the same, waves may have different wavelengths and frequencies.
Wavelength and frequency are inversely related.

11. Electromagnetic Spectrum
Electromagnetic spectrum – (EM spectrum) encompasses all forms of electromagnetic radiation, with the only differences in the types of radiation being their wavelengths and frequencies. (See Fig.5.5 on p.120)
Roy G Biv
Energy increase with frequency

12. Practice Problems
In your notes, complete the following practice problems as we work them together in class.
p.121 (1-4)

13. Particle Nature of Light
The concept of light as a wave explains much of its everyday behavior, but fails to describe aspects of light’s interactions with matter.
Why do heated objects emit only certain frequencies of light at a given temperature?
Why do some metals emit electrons when colored light of a specific frequency shines on them?

14. The Quantum Concept
Remember – the temperature of an object is a measure of the average kinetic energy of its particles… as temperature rises, particle energy increases.
For example – as iron is heated it possesses a greater amount of energy and emits different colors of light (gray – red – blue); The different colors correspond to different frequencies and wavelengths.

15. 1900 – German physicist Max Planck
studied the light emitted form heated objects. He concluded that matter can gain or lose energy only in small, specific amounts called quanta.
Quantum – the minimum amount of energy that can be gained or lost by an atom.
Equantum=hv (E=energy; h=Planck’s constant; and v=frequency)
Planck’s constant (h) – has a value of x Js
Energy of radiation increases as frequency increases.
For a given frequency – matter can emit or absorb energy only in whole-number multiples of hv; that is, 1hv, 2hv, 3hv,etc. Matter can only have certain amounts of energy – quantities of energy between these values do not exist.

16. The Photoelectric Effect
In the photoelectric effect, electrons, called photoelectrons, are emitted from a metal’s surface when light of a certain frequency shines on the surface.
Photoelectric cells convert the energy of incident light into electrical energy.
Ex. Solar calculators

17. Albert Einstein
proposed that electromagnetic radiation has both wavelike and particlelike natures. (the behavior of light is explained by a dual wave-particle model)
A beam of light has many wavelike characteristics, but can also be thought of as a stream of tiny particles, or bundles of energy, called photons.
Photon – a particle of electromagnetic radiation with no mass that carries a quantum of energy.
Ephoton=hv

18. Practice Problems
In your notes, complete the following practice problems as we work them together in class.
p.124 (5-6)

19. Atomic Emission Spectra
Atomic emission spectrum – the set of frequencies of the electromagnetic waves emitted by atoms of an element.
Each element’s atomic emission spectrum is unique and can be used to determine if that element is part of an unknown compound.

20. Flame Test Demonstration
Record your observations as Mrs. Orgeron demonstrates a flame test.

21. Assignment: Workbook p.25-26
Textbook p. 126 (even numbers only) Write the questions and answers!

22. Quantum Theory and the Atom
Chapter 5 Section 2:
Quantum Theory and the Atom

23. Questions still remained:
Reminder – the behavior of light is explained by a dual wave-particle model.
Questions still remained:
What is the relationship among atomic structure, electrons, and atomic emission spectra?

24. Bohr Model of the Atom
Niels Bohr (Danish physicist working with Rutherford) proposed a quantum model for the hydrogen atom that seemed to explain why atomic emission spectra were discontinuous.
Building on Planck’s and Einstein’s quantum theory, Bohr proposed that the hydrogen atom has only certain allowable energy states.
Ground state – lowest allowable energy state of an atom
When an atom gains energy, it is said to be in the excited state.
He suggested that the electron moves around the nucleus in only certain allowed circular orbits.
The smaller the orbit – the lower the atom’s energy state or energy level.
ΔE = Ehigher-energy orbit – Elower-energy orbit
Explains atomic emission spectra for Hydrogen, but no other elements.

25. The Quantum Mechanical Model of the Atom
1924 – Louis de Broglie (French grad student – physics) proposed an idea that accounted for fixed energy levels.
If waves can behave like particles, can particles behave like waves?
De Broglie equation – predicts that all moving particles have wave characteristics; λ= h/mv

26. The Heisenberg Uncertainty Principle
Werner Heisenburg (German theoretical physicist) concluded that it was impossible to make any measurement on an object without disturbing the object a little. (Think : finding a balloon in a dark room)
Heisenberg uncertainty principle – states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.
1926 – Austrian physicist Erwin Schrödinger derived an equation that treated the hydrogen atom’s electron as a wave; it also applied equally well to other elements.

27. Quantum Mechanical Model of the Atom –
the atomic model in which electrons are treated as waves (AKA the wave mechanical model)
Limits an electron’s energy to certain values
Makes no attempt to describe the electron’s path around the nucleus
Each solution to the equation is known as a wave function, which is related to the probability of finding an electron within a certain volume of space around the nucleus.
Atomic orbital – three dimensional region around the nucleus that describes the electron’s location

28. Hydrogen’s Atomic Orbitals
Principal quantum numbers – (n) indicates the relative sizes and energies of atomic orbitals.
As n increases, the orbital becomes larger, and the electron spends more time farther from the nucleus.
Therefore, n specifies the atom’s major energy levels, called principal energy levels.
Principal energy levels contain energy sublevels.
Level 1 = 1 sublevel
Level 2 = 2 sublevels
Level 3 = 3 sublevels
Etc.
Sublevels are labeled s, p, d, or f according to the shapes of an atom’s orbitals.
s – spherical (1 orbital)
p – dumbbell (3 orbitals)
d – not consistent (5 orbitals)
f – not consistent (7 orbitals)
Each orbital may contain 2 electrons.
Maximum number of orbitals related to each principal energy level is n2.
Maximum number of electrons related to each principle energy level is 2n2.

30. Electron Energy Levels

31. Assignments:
Workbook p.27-28

32. Electron Configurations
Chapter 5 Section 3:
Electron Configurations

33. Ground-State Electron Configurations
Electron configuration- the arrangement of electrons in an atom
Because low-energy systems are more stable than high-energy systems, electrons in an atom tend to assume the arrangement that gives the atom the lowest possible energy – the ground state electron configuration.
Three rules/principles define how electrons can be arranged in an atom’s orbitals.
The aufbau principle
The Pauli exclusion principle
Hund’s rule

34. The Aufbau Principle –
states that each electron occupies the lowest energy orbital available.
Aufbau diagram – (fig.5-17 p.135) the sequence of atomic orbitals from lowest energy to highest energy
All orbitals related to an energy sublevel are of equal energy.
In a multi-electron atom, the energy sublevels within a principal energy level have different energies.
In order of increasing energy, the sequence of energy sublevels within a principal energy level is s, p, d, and f.
Orbitals related to energy sublevels within one principal energy level can overlap orbitals related to energy sublevels within another principal energy level.

35. The Pauli exclusion principle
states that a maximum of two electrons may occupy a single atomic orbital, but only if the electrons have opposite spins.
Electrons are only able to spin in one of two directions: ↑ or ↓
An atomic orbital containing paired electrons with opposite spins is written as ↑↓.

36. Hund’s rule –
states that single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbital.

37. Orbital Diagrams & Electron Configuration Notations
Orbital diagrams – include a box for each of the atom’s orbitals
Example – Carbon
Electron configuration – shorthand method describing electron arrangement
Example – Carbon 1s2 2s2 2p2
Example - Sodium 1s2 2s2 2p6 3s1 or [Ne]3s1
See Table 5-3 p.137 and Table 5-4 p.138 for additional examples.
Figure 5-19 p.138 – the proper sequence for filling orbitals

40. Practice Problems:
In your notes, complete the following practice problems as we work them together in class.
p.139 (18-22)

41. Valence Electrons
Valence electrons are electrons in the outermost orbitals.
Determine the chemical properties of an element
On the periodic table –
Group 1 – 1 valence electron
Group 2 – 2 valence electrons
Group 13 – 3 valence electrons
Group 14 – 4 valence electrons
Group 15 – 5 valence electrons
Group 16 – 6 valence electrons
Group 17 – 7 valence electrons
Group 18 – 8 valence electrons (except He, which only has 2 e-)

42. Electron – Dot Structure
consists of the element’s symbol, which represents the atomic nucleus and inner-level electrons, surrounded by dots representing the atom’s valence electrons.
Example

43. Practice Problems:
In your notes, complete the following practice problems as we work them together in class.
p.141 (23)

44. Assignment:
Workbook p.29-30
Textbook p.141 (25-26, 28)