1. Chapter 5
Electrons in Atoms

2. Light and Quantized Energy
Section 1
Light and Quantized Energy

3. Section 1: Light and Quantized Energy
Light, a form of electromagnetic radiation, has characteristics of both a wave and a particle. K
What I Know W
What I Want to Find Out L
What I Learned

4. The Atom and Unanswered Questions
Recall that in Rutherford's model, the atom’s mass is concentrated in the nucleus and electrons move around it.
The model doesn’t explain how the electrons were arranged around the nucleus.
The model doesn’t explain why negatively charged electrons aren’t pulled into the positively charged nucleus.
In the early 1900s, scientists observed certain elements emitted visible light when heated in a flame.
Analysis of the emitted light revealed that an element’s chemical behavior is related to the arrangement of the electrons in its atoms.
Light and Quantized Energy

5. Flame test

6. The Wave Nature of Light
Visible light is a type of electromagnetic radiation, a form of energy that exhibits wave-like behavior as it travels through space.
All waves can be described by several characteristics.
The wavelength (λ) is the shortest distance between equivalent points on a continuous wave.
The frequency (ν) is the number of waves that pass a given point per second.
The amplitude is the wave’s height from the origin to a crest.
Light and Quantized Energy

7. The Wave Nature of Light
The speed of light (3.00 × 108 m/s) is the product of it’s wavelength and frequency c = λν.
Light and Quantized Energy

8. The Wave Nature of Light
Sunlight contains a continuous range of wavelengths and frequencies.
A prism separates sunlight into a continuous spectrum of colors.
The electromagnetic spectrum includes all forms of electromagnetic radiation.
Light and Quantized Energy

9. Electromagnetic Spectrum

10. Properties of Waves Wave length: The length of each wave
λ = c/v
Constant: speed of light (c)
3.00 × 108 m/s
Frequency (v) Hz
Wavelength (λ) m c λ v

11. Calculations
λ = c/v
Microwaves are used to cook food and transmit information. What is the wavelength of a microwave that has a frequency of × 109 Hz?
KNOWN
UNKNOWN
ν = 3.44 × 109 Hz
λ = ? m
c = 3.00 × 108 m/s

12. Calculations
λ = c/v
What is the frequency of a microwave that has a wavelength of 6.72 × 10-9 m?
KNOWN
UNKNOWN
λ =
ν =
c =

13. The Particle Nature of Light
The wave model of light cannot explain all of light’s characteristics. Some examples include:
Why heated objects emit only certain frequencies of light at a given temperature.
Why some metals emit electrons when light of a specific frequency shines on them.
Light and Quantized Energy

14. The Particle Nature of Light
The Quantum Concept
In 1900, German physicist Max Planck ( ) began searching for an explanation of this phenomenon as he studied the light emitted by heated objects.
Planck’s study led him to a startling conclusion:
Matter can gain or lose energy only in small, specific amounts called quanta.
A quantum is the minimum amount of energy that can be gained or lost by an atom.
Planck’s constant has a value of × 10–34 J ● s.
Light and Quantized Energy

15. The Particle Nature of Light
The Photoelectric Effect
The photoelectric effect is when electrons are emitted from a metal’s surface when light of a certain frequency shines on it.
Light and Quantized Energy

16. Photoelectric Effect

17. The Particle Nature of Light
Light’s Dual Nature
Albert Einstein proposed in 1905 that light has a dual nature.
A beam of light has wavelike and particle-like properties.
A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.
Light and Quantized Energy

18. Planck

19. Energy of a Photon Ephoton energy Frequency (v) Hz
Ephoton = hv
Ephoton
energy
Frequency (v) Hz
H (Planck’s contstant)
6.626 × 10–34 J ● s.
Ephoton h v

20. Calculations
Every object gets its color by reflecting a certain portion of incident light. The color is determined by the wavelength of the reflected photons, thus by their energy. What is the energy of a photon from the violet portion of the Sun’s light if it has a frequency of × 1014 s-1?
Ephoton = hv
KNOWN
UNKNOWN
ν = × 1014 s-1
Ephoton = ? J
h = × J•s

21. Atomic Emission Spectra
Light in a neon sign is produced when electricity is passed through a tube filled with neon gas and excites the neon atoms.
The excited atoms return to their stable state by emitting light to release energy.
Light and Quantized Energy

22. Atomic Emission Spectra
The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by the atoms of the element.
Each element’s atomic emission spectrum is unique.
Light and Quantized Energy

23. Essential Questions
How do the wave and particle natures of light compare?
How is a quantum of energy related to an energy change of matter?
How do continuous electromagnetic spectra and atomic emission spectra compare and contrast?
Light and Quantized Energy

24. Quantum Theory and the Atom
Section 2
Quantum Theory and the Atom

25. Section 2: Quantum Theory and the Atom
Wavelike properties of electrons help relate atomic emission spectra, energy states of atoms, and atomic orbitals. K
What I Know W
What I Want to Find Out L
What I Learned

26. Bohr’s Model of the Atom
Einstein’s theory of light’s dual nature accounted for several unexplainable phenomena but not why atomic emission spectra of elements were discontinuous rather continuous.
In 1913, Niels Bohr, a Danish physicist working in Rutherford’s laboratory, proposed a quantum model for the hydrogen atom that seemed to answer this question.
Bohr correctly predicted the frequency lines in hydrogen’s atomic emission spectrum.
The lowest allowable energy state of an atom is called its ground state.
When an atom gains energy, it is in an excited state.
Quantum Theory and the Atom

27. Bohr’s Model of the Atom
Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits.
Quantum Theory and the Atom

28. Bohr’s Model of the Atom
Each orbit was given a number, called the quantum number.
Hydrogen’s single electron is in the n = 1 orbit in the ground state.
When energy is added, the electron moves to the n = 2 orbit.
Quantum Theory and the Atom

29. Atomic orbitals
Orbital- is a region of space around the nucleus where an electron is likely to be found.
An electron cloud is a good approximation of how electrons behave in their orbitals
The level in which an electron has the least energy—the lowest energy level—has only one orbital. Higher energy levels have more than one orbital

30. Electron Configuration
Energy Levels, Orbitals, and Electrons
The most stable electron configuration is the one in which the electrons are in orbitals with the lowest possible energies.
Energy Level
Number of Orbitals
Max. # of Electrons 1 2 4 8 3 9 18 16 32

31. DRAWING BOHR MODELS OF ATOMS
Draw the nucleus and label with #p & #n
Draw electron orbitals
1st orbital can have 2 electrons only
2nd and 3rd ring can each have _8___electrons
Fill – _First ring FIRST______
Inner rings must be filled first before any electron enters a higher ring!!!!!
Let’s Practice:

32. Bohr Model http://www. youtube. com/watch

33. Bozeman Science: Bohr

34. Bohr Model - Hydrogen

35. Bohr Model - Oxygen

36. Bohr Model - Sulfur

37. Bohr’s Model of the Atom
The limits of Bohr’s model
Bohr’s model explained the hydrogen’s spectral lines, but failed to explain any other element’s lines.
The behavior of electrons is still not fully understood, but substantial evidence indicates they do not move around the nucleus in circular orbits.
Quantum Theory and the Atom

38. The Quantum Mechanical Model of the Atom
Heisenberg showed it is impossible to take any measurement of an object without disturbing it.
The Heisenberg uncertainty principle states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.
The only quantity that can be known is the probability for an electron to occupy a certain region around the nucleus.
Quantum Theory and the Atom

39. The Quantum Mechanical Model of the Atom
Schrödinger treated electrons as waves in a model called the quantum mechanical model of the atom.
Schrödinger’s equation applied equally well to elements other than hydrogen.
The wave function predicts a three-dimensional region around the nucleus called the atomic orbital.
Quantum Theory and the Atom

40. Hydrogen Atomic Orbitals
Principal quantum number (n) indicates the relative size and energy of atomic orbitals.
n specifies the atom’s major energy levels, called the principal energy levels.
Energy sublevels are contained within the principal energy levels.
Quantum Theory and the Atom

41. Hydrogen Atomic Orbitals
Each energy sublevel relates to orbitals of different shape.
Quantum Theory and the Atom

42. Hydrogen Atomic Orbitals
Quantum Theory and the Atom

43. Essential Questions
How do the Bohr and quantum mechanical models of the atom compare?
What is the impact of de Broglie’s wave-particle duality and the Heisenberg uncertainty principle on the current view of electrons in atoms?
What are the relationships among a hydrogen atom’s energy levels, sublevels, and atomic orbitals?
Quantum Theory and the Atom

44. Electron Configuration
Section 3
Electron Configuration

45. Section 3: Electron Configuration
Three rules are used to determine electron arrangement in an atom. K
What I Know W
What I Want to Find Out L
What I Learned

46. Bozeman Science: Electron Configuration

47. Ground-State Electron Configuration
The arrangement of electrons in the atom is called the electron configuration.
The aufbau principle states that each electron occupies the lowest energy orbital available.
Electron Configuration

48. Ground-State Electron Configuration

49. Ground-State Electron Configuration
The aufbau diagram can be used to write correct ground-state electron configurations for all elements up to and including Vanadium, atomic number 23.
The electron configurations for certain transition metals, like chromium and copper, do not follow the aufbau diagram due to increased stability of half- filled and filled sets of s and d orbitals.
Electron Configuration

50. Electron Configuration
Electron Configuration - a representation of the arrangement of electrons in an atom

51. Electron Configuration
Examples of electron Configuration
1. Li 1s22s1
2. C 1s22s22p6
# of e- in that shell
principle
azimuthal

52. Electron Configuration
Take note that after 4s is filled, 3d is than filled before 4p.
…… 6s than 4f than 5d than 6p
When writing out the electron configuration, always write your numbers in numerical order
Y 1s22s22p63s23p64s23d104p65s24d1 – NO!
Y 1s22s22p63s23p63d104s24p64d15s2

53. Electron Configuration
Examples: Be O Ca Mn

54. Electron Configuration
Examples Pb Os

55. Ground-State Electron Configuration
Hund’s Rule (better known as the Bus Rule)
Before any second electron can be placed in a sub level, all the orbitals of that sub level must contain at least one electron – spread out the e- before pairing them up.
Pauli Exclusion Principle - electrons occupying the same orbital must have opposite spin.
See a good online illustration at

56. Orbital Notation

57. Orbital Notation H 1s F 1s 2s 2p

58. Orbital Notation
Examples:
Li F
Na Sc

59. Valence Electrons
Valence electrons are defined as electrons in the atom’s outermost orbitals—those associated with the atom’s highest principal energy level.
Electron-dot structure consists of the element’s symbol representing the nucleus, surrounded by dots representing the element’s valence electrons.
Electron Configuration

60. Significance of Electron Configurations
Valence shell electrons - outermost electrons involved with bonding
no atom has more than 8 valence electrons
Noble gases - 8 valence electrons – least reactive of all elements
Lewis Dot structures: NSEW (cheating) also show correct way, count to 8
Lewis Dot Structures

61. Essential Questions
How are the Pauli exclusion principle, the aufbau principle, and Hund’s rule used to write electron configurations?
How do electron-dot structures represent an atom’s valence electrons?
Electron Configuration