1. Chapter 5 - Electrons in Atoms

2. THE NUCLEAR ATOM AND UNANSWERED QUESTIONS
1. In the early 1900’s scientists were able to
determine that an element’s chemical behavior is related to the arrangement of electrons in its atoms

3. WAVE NATURE OF LIGHT
1. Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space
Examples: a) visible light (ROYGBIV)
b) microwaves
c) x-rays
d) radio waves

4. WAVE NATURE OF LIGHT
2. Primary characteristics of all waves include wavelength, frequency, amplitude, and speed
SEE Fig. 5-2, Page 119
3. Wavelength () is the shortest distance
between equivalent points on a continuous
wave is usually expressed in meters,
centimeters, or nanometers

5. WAVE NATURE OF LIGHT
4. Frequency (  ) is the number of waves that pass a given point per second
5. The SI unit for frequency is the hertz (Hz)
which equals one wave per second
Example: 652 Hz = 652 waves/second =
652/s = 652s-1

6. WAVE NATURE OF LIGHT
6. The amplitude of a wave is the wave height from the origin to a crest, or from the origin to a trough
7. All electromagnetic waves travel at the speed of 3.00 X 108 m/s in a vacuum

7. WAVE NATURE OF LIGHT
8. The speed of light ( c ) is the product of its
wavelength (  ) and its frequency (  )
c =  
9. Although the speed of all electromagnetic
waves is the same, waves may have different wavelengths and frequencies
SEE Fig. 5-3, Page 119

8. WAVE NATURE OF LIGHT
10. White light is a combination of all the colors of the visible spectrum
R O Y G B I V
e r e r l n i
d a l e u d o
n l e e i l
g o n g e
e w o t
SEE Fig. 5-4, Page 1`20

9. WAVE NATURE OF LIGHT
11. The electromagnetic spectrum encompasses all forms of electromagnetic radiation
SEE Fig. 5-5, Page 120
Example Problem 5-1, Page 121
Practice Problems, Page 121

10. Practice Problems
1. A helium-neon laser emits light with a
wavelength of 633 nm. What is the frequency of this light?
2. What is the wavelength of X-rays having a frequency of 4.80 X 1017 Hz?
3. An FM radio station broadcasts at a frequency of 98.5 MHz. What is the wavelength of the station’s broadcast signal?

11. PARTICLE NATURE OF LIGHT
1. The temperature of an object is a measure
of the average kinetic energy of its particles
2. As an object gets hotter it possesses a greater amount of energy, and emits different colors of light. SEE Fig. 5-6, Page 122

12. PARTICLE NATURE OF LIGHT
3. These different colors correspond to different frequencies and wavelengths
4. Max Planck determined that matter can gain or lose energy only in small, specific amounts called quanta

13. PARTICLE NATURE OF LIGHT
5. A quantum is the minimum amount of energy that can be gained or lost by an atom
6. The energy of a quantum is related to the
frequency of the emitted radiation by the
equation
Equantum = h 
where E is energy, h is Planck’s constant, and  is frequency

14. PARTICLE NATURE OF LIGHT
7. Planck’s constant has a value of
6.626 X J•s
8. The photoelectric effect is a phenomenon in
which photoelectrons are emitted from a
metal’s surface when light of a certain
frequency shines on the surface
SEE Fig. 5-7, Page 123
Example: Calculator in Fig. 5-8, Page 123

15. PARTICLE NATURE OF LIGHT
9. While a beam of light has many wave
characteristics, it can also be thought of as a stream of tiny particles, or bundles of energy called photons
10. A photon is a particle of electromagnetic
radiation with no mass that carries a quantum of energy

16. PARTICLE NATURE OF LIGHT
11. Albert Einstein calculated that a photon’s energy depends on its frequency
Ephoton = h 
Example Problem 5-2, Page 124
Practice Problems, Page 124

17. ATOMIC EMISSION SPECTRA
1. The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by atoms of the element
2. The atomic emission spectrum is characteristic of the element and is used to identify the element.
Section 5.1 Assessment, Page 126

18. BOHR MODEL OF THE ATOM
1. Niels Bohr proposed a quantum model for atoms which included:
a) predicted the frequencies of the lines in
hydrogen’s atomic emissions spectrum
b) proposed that the hydrogen atom has only certain allowable energy states

19. BOHR MODEL OF THE ATOM
c) related the hydrogen atom’s energy states to the motion of the electron within the atom
d) the single electron in a hydrogen atom
moves around the nucleus in only certain allowed circular orbits

20. e) the smaller an electron’s orbit, the lower
BOHR MODEL OF THE ATOM
e) the smaller an electron’s orbit, the lower
the atom’s energy state, or energy level
f) assigned a quantum number, (n), to each
orbit and calculated the orbits radius
SEE Table 5-1, Page 127

21. BOHR MODEL OF THE ATOM
2. The lowest allowable energy state of an atom is called its ground state
3. In the ground state, an atom does not radiate energy
4. When energy is added to an atom by an
outside force, electrons can move up to a
higher-energy orbit (called the excited state)

22. BOHR MODEL OF THE ATOM
5. Electrons can also drop to a lower-energy level and emit energy in the form of photons
6. To calculate the energy (E) of a photon we can use the following equation:
ΔE = Ehigher-energy orbit - Elower-energy orbit = Ephoton = h 

23. THE QUANTUM MECHANICAL MODEL OF THE ATOM
1. Louis de Broglie proposed that particles of matter can behave like waves
2. de Broglie derived an equation for the
wavelength (  ) of a particle of mass (m)
moving at a velocity ( v )
 = h mv

24. 3. The de Broglie equation predicts that all
moving particles have wave characteristics
Example: 910 kg traveling at 25 m/s
m = 910 kg
v = 25 m/s
h = X J
 = ?
 = h = X J = 2.9 X m
mv (910kg)(25m/s)

25. 1. Werner Heisenberg concluded that it is
THE HEISENBERG UNCERTAINTY PRINCIPLE
1. Werner Heisenberg concluded that it is
impossible to make any measurement on an
object without disturbing the object
2. Consider the energy of a photon:
A high-energy photon of electromagnetic
radiation has about the same energy as an
electron.
The interaction between the two particles
changes both the wavelength of the photon
and the position and velocity of the electron
SEE Fig. 5-13, Page 132

26. 3. 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.
Example: The effect of a photon emitted by a
flashlight on a helium balloon is so
small that it is virtually impossible
to measure
4. Erwin Schrodinger derived an equation that treated the hydrogen atom’s electron as a wave

27. 6. The region around the nucleus of an atom
THE HEISENBERG UNCERTAINTY PRINCIPLE
5. The atomic model proposed by Schrodinger is called the quantum mechanical model of the atom
6. The region around the nucleus of an atom
where electrons are located is called an atomic orbital
SEE Fig. 5-13, Page 132

28. HYDROGEN’S ATOMIC ORBITALS
1. Principle Quantum Numbers ( n ) indicate
the relative sizes and energies of atomic
orbitals
2. As n increases, the orbital becomes larger, the electron spends more time farther from the nucleus, and the atom’s energy level increases

29. HYDROGEN’S ATOMIC ORBITALS
3. n specifies the atom’s major energy levels,
called principle energy levels
4. Up to seven energy levels have been detected
5. Principle energy levels contain energy
sublevels

30. HYDROGEN’S ATOMIC ORBITALS
6. The quantum number assigned to an energy level also indicates the number of sublevels contained within the energy level
Examples:
First energy level (n = 1) has one sublevel
Second energy level (n = 2) has two sublevels
Third energy level (n = 3) has three sublevels

31. n = 3 has a 3s, 3p, and 3d sublevel
HYDROGEN’S ATOMIC ORBITALS
7. Sublevels are labeled s, p, d, or f according to the shape of the atoms orbitals
Examples:
n = 1 has a 1s sublevel
n = 2 has a 2s and 2p sublevel
n = 3 has a 3s, 3p, and 3d sublevel
n = 4 has a 4s, 4p, 4d, and 4f sublevel
SEE Table 5-2, Page 134
Section 5.2 Assessment, Page 134

32. 1. The arrangement of electrons in an atom is
GROUND STATE ELECTRON CONFIGURATION
1. The arrangement of electrons in an atom is
called the atom’s electron configuration
2. Electrons in an atom tend to assume the
arrangement that gives the atom the lowest
possible energy
3. The most stable, lowest-energy arrangement of the electrons in an atom of each element is called the element’s ground-state electron configuration

33. 4. The aufbau principle states that each
GROUND STATE ELECTRON CONFIGURATION
4. The aufbau principle states that each
electron occupies the lowest energy orbital
available
5. The sequence of atomic orbitals from lowest to highest energy (aufbau diagram) is shown in Figure 5-17, Page 135

34. Aufbau Diagram 7p

35. 6. Key features of the aufbau principle include:
GROUND STATE ELECTRON CONFIGURATION
6. Key features of the aufbau principle include:
a) all orbitals related to an energy sublevel
are of equal energy
b) in a multi-electron atom, energy sublevels within a principal energy level have different energies

36. GROUND STATE ELECTRON CONFIGURATION
c) in order of increasing energy, the sequence of energy sublevels within a principal energy level is s, p, d, and f
d) orbitals related to energy sublevels within one principal energy level can overlap orbitals related to energy sublevels within another principal level

37. The Pauli Exclusion Principle
7. The Pauli Exclusion principle states that a maximum of two electrons may occupy a single orbital, and the electrons must have opposite spins
Each type of subshell contains a different number of orbitals. 
Each orbital can hold at most 2 electrons. 

38. The Pauli Exclusion Principle
The following table shows how many electrons each type of subshell can hold.
Subshell Type
# of Orbitals
Maximum # of Electrons s 1 2 p 3 6 d 5 10 f 7 14

39. Table 3-6b Orbitals and Electron Capacity of the First Four Principle Energy Levels
Principle energy level (n)
Type of sublevel
Number of orbitals per type
Number of orbitals per level(n2)
Maximum number of electrons (2n2) 1 s 2 4 8 p 3 9 18 d 5 16 32 f 7   

40. 8. The Pauli Exclusion principle states that a
GROUND STATE ELECTRON CONFIGURATION
7. Electrons can spin in one of two directions. () represents one direction and () represents the opposite direction
8. The Pauli Exclusion principle states that a
maximum of two electrons may occupy a
single orbital, and the electrons must have
opposite spins

41. GROUND STATE ELECTRON CONFIGURATION
9. An atomic orbital containing paired electrons with opposite spin are written as 
10. 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 orbitals

42. ORBITAL DIAGRAMS AND ELECTRON
CONFIGURATION NOTATIONS
1. Two methods of representing an atom’s
electron configuration are:
a) orbital diagram
b) electron configuration notation
SEE Table 5-3, Page 137

43. 2. Using Sodium (Na) as an example:
Na 1s22s22p63s1
     
1s 2s p s
3. Electron configurations for noble gases use
bracketed symbols
Example: a) Helium = [ He]
b) Neon = [ Ne ]

45. ORBITAL DIAGRAMS AND ELECTRON
CONFIGURATION NOTATIONS
4. The electron configuration for an element can be represented using the noble-gas notation for the noble gas in the previous period and the electron configuration for the energy level
being filled
Example: Sodium (Na) = [Ne]3s1

46. Topic for 11.03.05 Ch. 5 Electrons in Atoms
Discuss Valence electrons and Electron-dot structure

47. ORBITAL DIAGRAMS AND ELECTRON CONFIGURATION NOTATIONS
5. SEE Table 5-4 for electron configurations for elements in period three
6. A memory aid called a sublevel diagram is
shown in Fig. 5-19, Page 138
Example Problem 5-3, Page 139
Sample Problems, Page 139

48. VALENCE ELECTRONS
1. Valence electrons are defined as electrons in the atom’s outermost orbital
2. Number of valence electrons is the same as an elements group number (A-families)

49. VALENCE ELECTRONS
3. Number of energy levels is the same as an
elements period number
4. An atom’s electron-dot structure consists of the elements symbol surrounded by dots representing each of the atom’s valence electrons
SEE Table 5-5, Page 140

50. Ch 5 test preparation:
Example Problem 5-4, Page 141
Practice Problems, Page 141
Section 5.3 Assessment, Page 141
Chapter 5 Assessment, Pages
Standardized Test Practice - Ch. 5, Page 149