1. HONORS CHEMISTRY Atomic Structure and Electrons

2. Light and Quantized Energy
Chemists found Rutherford’s nuclear model lacking because it did not begin to account for the differences in chemical behavior among the various elements.
In the early 1900s, scientists observed that certain elements emitted visible light when heated in a flame.

3. Light and Quantized Energy
Analysis of the emitted light revealed that an element’s chemical behavior is related to the arrangement of the electrons in its atoms.
In order to better understand this relationship and the nature of atomic structure, it will be helpful to first understand the nature of light.

4. The Electromagnetic Spectrum
Electromagnetic radiation includes radio waves that carry broadcasts to your radio and TV, microwave radiation used to heat food in a microwave oven, radiant heat used to toast bread, and the most familiar form, visible light.
All of these forms of radiant energy are parts of a whole range of electromagnetic radiation called the electromagnetic spectrum.

5. The Electromagnetic Spectrum

6. Particle Nature of Light
While considering light as a wave does explain much of its everyday behavior, it fails to adequately describe important aspects of light’s interactions with matter.

7. The Quantum Concept
In 1900, the German physicist Max Planck (1858–1947) began searching for an explanation as he studied the light emitted from heated objects.

8. The Quantum Concept
His study of the phenomenon led him to a startling conclusion: matter can gain or lose energy only in small, specific amounts called quanta.
That is, a quantum is the minimum amount of energy that can be gained or lost by an atom.

9. The Quantum Concept
While a beam of light has many wavelike characteristics, it also can be thought of as a stream of tiny particles, or bundles of energy, called photons
Thus, a photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.

10. The Heisenberg Uncertainty Principle
Heisenberg concluded that it is impossible to make any measurement on an object without disturbing the object—at least a little.
The act of observing the electron produces a significant, unavoidable uncertainty in the position and motion of the electron.

11. The Heisenberg Uncertainty Principle
Heisenberg’s analysis of interactions such as those between photons and electrons led him to his historic conclusion.
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.

12. Combining them : E=hc/l
Mathematic Equations
c=ln
c= speed of light (3.0 x 108m/s)
l= wavelength (m)
n= frequency (s-1 or Hz)
E=hn
E=energy (J or kg∙m2/s2)
h=Planck’s constant (6.63x10-34 J∙s or kg∙m2/s)
n=frequency (s-1)
Combining them : E=hc/l

13. Energy levels, sublevels, & orbitals
Energy levels=clouds or shells around nucleus (n=1,2,3…)
Sublevels=found inside energy levels (s,p,d,f)
Atomic orbitals=found within sublevels:
s = 1 orbital (sphere)
p = 3 orbitals (dumbell)
d = 5 orbitals (p. 313)
f = 7 orbitals
2 e- max per orbital

14. Rules Governing e- Configurations
Aufbau Principle = e- fill orbitals with lowest energy first
Pauli Exclusion Principle = e- in the same orbital have opposite spins → no 2 e- in a single atom will have the same set of quantum numbers
Hund’s Rule = e- occupy one orbital in each sublevel before pairing up (p,d,f)

15. Electron Configurations
Use Periodic Table to find e- configurations

16. Electrons
Diamagnetism = all of e- are paired; not strongly affected by magnetic fields
Paramagnetism = has unpaired e-; strongly affected by magnetic fields
Valence e- = e- in outermost energy level
For Representative Elements 1A-8A, groups number=number of valence e-
Period number=energy level of valence e-

17. Quantum Numbers