1. Chapter 5 Periodicity and the Electronic Structure of Atoms
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Chapter 5 Periodicity and the Electronic Structure of Atoms

2. A Theory that Explains Electron Behavior
In everyday language, small is a relative term: something is smaller than something else.
Atoms and the particles that compose them are unimaginable small, electron has the mass of less than a trillionth of a trillion of a gram and a size that is immeasurable. And yet, an atom’s electrons, determine many of its chemical and physical properties.
The quantum-mechanical model explains the manner electrons exist and behave in atoms
helps us understand and predict the properties of atoms that are directly related to the behavior of the electrons
why some elements are metals while others are nonmetals
why some elements gain 1 electron when forming an anion, while others gain 2
why some elements are very reactive while others are practically inert
and other Periodic patterns we see in the properties of the elements

3. The Nature of Light its Wave Nature
One of the ways that energy travels through space:
– Light from sun; microwave oven; radio waves for MRI mapping
• Exhibit the same type of wavelike behavior and
travel at the speed of light in a vacuum
• It has electric and magnetic fields that
simultaneously oscillate in planes mutually
perpendicular to each other and to the direction
of propagation through space.
Electromagnetic radiation has oscillating electric (E) and magnetic (H) fields in planes

4. The Nature of Light its Wave Nature
light is a form of electromagnetic radiation
composed of perpendicular oscillating waves, one for the electric field and one for the magnetic field
an electric field is a region where an electrically charged particle experiences a force
a magnetic field is a region where an magnetized particle experiences a force
all electromagnetic waves move through space at the same, constant speed
3.00 x 108 m/s in a vacuum = the speed of light, c

5. The Electromagnetic Spectrum
visible light comprises only a small fraction of all the wavelengths of light – called the electromagnetic spectrum
short wavelength (high frequency) light has high energy
radiowave light has the lowest energy
gamma ray light has the highest energy
high energy electromagnetic radiation can potentially damage biological molecules
ionizing radiation l 

6. Characterizing Waves the number of waves = number of cycles
Waves are characterized by wavelength, frequency, and speed.
– wavelength (λ) is the distance between two
consecutive peaks or troughs in a wave.
– frequency (ν) is defined as the number of waves (cycles) per second
- is a measure of the distance covered by the wave
the distance from one crest to the next
the number of waves = number of cycles
units are hertz, (Hz) or cycles/s = s-1
1 Hz = 1 s-1
the amplitude is the height of the wave
the distance from node to crest
or node to trough
the amplitude is a measure of how intense the light is – the larger the amplitude, the brighter the light

7. Amplitude & Wavelength
Dim light
Bright light

8. The Relationship Between Wavelength and Frequency
for waves traveling at the same speed, the shorter the wavelength, the more frequently they pass
this means that the wavelength and frequency of electromagnetic waves are inversely proportional
since the speed of light is constant, if we know wavelength we can find the frequency, and visa versa
c = v x λ
c is defined to be the rate of travel of all electromagnetic energy in a vacuum and is a constant value—speed of light.
c= 3.00 x 108 m/s

9. Example
Calculate the wavelength of red light with a frequency of 4.62 x 1014 s-1
A laser dazzles the audience in a rock concert by emitting green light with a wave length of 515 nm. Calculate the frequency of the light

10. Particlelike Properties of Radiant Energy: The Photoelectric Effect and Planck’s Postulate
it was observed that many metals emit electrons when a light shines on their surface
this is called the Photoelectric Effect
classic wave theory attributed this effect to the light energy being transferred to the electron
according to this theory, if the wavelength of light is made shorter, or the light waves intensity made brighter, more electrons should be ejected

11. Einstein’s Explanation
In 1905, Einstein proposed a bold explanation: light energy must come in packets, called quanta or photons
Energy is in fact quantized and can be transferred only in discrete units of size hν.
the energy of a photon of light was directly proportional to its frequency
inversely proportional to it wavelength
the proportionality constant is called Planck’s Constant, (h) and has the value x J∙s
1 Joules = kg•m2/s2

12. Examples
What is the energy (in kJ/mol) of photons of radar waves with ν = 3.35 x 108 Hz?
Calculate the number of photons in a laser pulse with wavelength 337 nm and total energy 3.83 mJ

13. The Bohr Model of the Atom: Quantized Energy
In 1912, Neils Bohr proposed that the electrons could only have very specific amounts of energy
fixed amounts = quantized
the electrons traveled in orbits that were a fixed distance from the nucleus
stationary states
therefore the energy of the electron was proportional the distance the orbital was from the nucleus
electrons emitted radiation when they “jumped” from an orbit with higher energy down to an orbit with lower energy
the distance between the orbits determined the energy of the photon of light produced

14. The Bohr Model of the Atom: Quantized Energy
In the Bohr model, each spectral line is produced when an electron falls from one stable orbit, or stationary state, to another lower energy.
Although his model only works well for one-electron atoms like Hydrogen, it fails for multiple-electron atoms
The modern model of electron has discarded “orbits” but retained the idea of energy of electron is “quantized”.

15. Wavelike Properties of Matter
Chapter 5: Periodicity and Atomic Structure
Wavelike Properties of Matter
2/19/2019
Energy is really a form of matter, and all matter exhibits both particulate and wave properties.
–Large “pieces” of matter, such as base balls, exhibit predominantly particulate properties
–Very small “pieces” of matter, such as photon, while showing some particulate properties through relativistic effects, exhibit dominantly wave properties
–“Pieces” of intermediate mass, such as electrons, show both the particulate and wave properties of matter
Louis de Broglie in 1924 suggested that, if light can behave in some respects like matter, then perhaps matter can behave in some respects like light.
In other words, perhaps matter is wavelike as well as particlelike. mv h
l =
Copyright © 2008 Pearson Prentice Hall, Inc.

16. examples
What velocity would an electron (mass = 9.11 x 10-31kg) need for its de Broglie wavelength to be that of red light (750 nm)?
Determine the wavelength of a neutron traveling at 1.00 x 102 m/s (Massneutron = x g)

17. Quantum Mechanics and the Heisenberg Uncertainty Principle
As we just saw in the de Broglie relation, the velocity of an electron is related to its wave nature. The position of an electron, however, is related to its particle nature. ( Particles have well-defined position, but wave doe not).
Heisenberg Uncertainty Principle – both the position (Δx) and the momentum (Δmv) of an electron cannot be known beyond a certain level of precision
1. (Δx) (Δmv) > h 4π
2. Cannot know both the position and the momentum of an electron with a high degree of certainty
3. If the momentum is known with a high degree of certainty
i. Δmv is small
ii. Δ x (position of the electron) is large
4. If the exact position of the electron is known
i. Δmv is large
ii. Δ x (position of the electron) is small

18. The Quantum Mechanical Model of the Atom: Orbitals and Quantum Numbers
In 1926, Erwin Schrodinger proposed the quantum mechanical model of an atom, which is the modern description of how electrons exist around an atom. Quantum mechanic uses the dual nature (particle-like and wave-like) of matter to understand the behavior of electrons.
An orbit for an electron is thus forbidden because it would allow us to simultaneously know the exact position and velocity of electron, just like we can for satellites.
Violate Heisenberg Uncertainty principle
Instead, quantum mechanical focus on the wave-like nature of electron and proposes that electrons exist around nucleus as “wave-function” which are now known as orbitals.

19. The Quantum Mechanical Model of the Atom: Orbitals and Quantum Numbers
Solve
Wave
equation
Wave function
or orbital (Y)
Probability of finding
electron in a region
of space (Y 2)
calculations show that the size, shape and orientation in space of an orbital are determined be three integer terms in the wave function
added to quantize the energy of the electron
these integers are called quantum numbers
principal quantum number, n
angular momentum quantum number, l
magnetic quantum number, ml
Since we can’t ever be certain of the electron’s position, we work with probabilities.

20. The Quantum Mechanical Model of the Atom: Orbitals and Quantum Numbers
Principal Quantum Number (n)
Describes the size and energy level of the orbital
Commonly called a shell
Positive integer (n = 1, 2, 3, 4, …)
As the value of n increases,
the energy increases.
the average distance of the e– from the nucleus increases.

21. The Quantum Mechanical Model of the Atom: Orbitals and Quantum Numbers
Angular-Momentum Quantum Number (l)
Defines the three-dimensional shape of the orbital
Commonly called a subshell
There are n different shapes for orbitals.
If n = 1, then l = 0.
If n = 2, then l = 0 or 1.
If n = 3, then l = 0, 1, or 2.
Commonly referred to by letter (subshell notation)
l = 0 s (sharp)
l = 1 p (principal)
l = 2 d (diffuse)
l = 3 f (fundamental)
Scientists working in spectroscopy first used sharp, principal, diffuse, and fundamental with respect to atomic spectra.
After f, the series goes alphabetically (g, h, etc.).
the energy of the subshell increases with l (s < p < d < f).

22. The Quantum Mechanical Model of the Atom: Orbitals and Quantum Numbers
Magnetic Quantum Number (ml )
Defines the spatial orientation of the orbital
There are 2l + 1 values of ml, and they can have any integral value from -l to +l.
If l = 0, then ml = 0.
If l = 1, then ml = –1, 0, or 1.
If l = 2, then ml = –2, –1, 0, 1, or 2.

23. The Quantum Mechanical Model of the Atom: Orbitals and Quantum Numbers

24. Examples
Give the possible combination of quantum numbers for the following orbitals
3s orbital 2 p orbitals
Give orbital notations for electrons in orbitals with the following quantum numbers:
n = 2, l = 1 ml = 1 n = 3, l = 2, ml = -1

25. The Shapes of Orbitals _ memorize
Node: A surface of zero probability for finding the electron
s orbitals are spherical and penetrate closer to the nucleus.

26. The Shapes of Orbitals-memorize
p orbitals are dumbbell-shaped, with a nodal plane running through the nucleus.
Introductory Physical Chemistry courses often explore the mathematical relationships that define the regions of space (or shape).