1. Atoms and Periodic Properties
Chapter 8
Atoms and Periodic Properties
Good APCs: 1, 3, 4, 5, 6, 7, 9, 10, 11, 12, 13, 15, 18, 19, 23, 24, 25, 26, 29, 31, 32, 33, 34, 39, 40, 42, 43, 45, 46, 47, 49, 50
This shows the conceptual, computer-generated, model of a Beryllium Atom (Fig 8.1)
Nucleus and 1s, 2s, electron orbitals (which can be predicted using the Periodic Table)

2. Discovery of the Electron
J. J. Thomson (late 1800’s)
Fig 8.2
Fig 8.3
Performed cathode ray experiments (Figures 8.2 and 8.3)
Cathode Ray passed between two charged plates is deflected (animation on homepage)
Towards the positive charged plate by a magnetic field
FYI: Detecting screen was zinc sulfide
Visible light is produced when charged particle strikes it
Shows up as a greenish beam
Identified electron as a fundamental particle
Discovered negatively charged electron
Measured electron’s charge-to-mass ratio
Mass of the electron (at rest): x kg (see front of book)

3. Electron charge Robert Millikan (~1906) Fig 8.4
Studied charged oil droplets in an electric field (Figure 8.4)
Charge on droplets = multiples of electron charge
Charge + Thomson’s result gave electron mass
FYI: He had to follow an individual droplet and measure how fast it moved
Again, the mass of the electron is pretty darn small: 1/1840 of the mass of Hydrogen
Fig 8.4

4. The Nucleus Ernest Rutherford (1907) Fig 8.6 Fig 8.5 Good animation
Scattering of “alpha” particles off gold foil (page 226)
Radioactive matter producing an alpha particle
An alpha particle is a Helium atom that lost 2 electrons thus is positively charge (He+2)
Conclusion: an atom’s positive charge resides in a small, massive nucleus
“Solar System” Model with a lot of space
Later: positive charges = protons
James Chadwick (1932): also neutral neutrons in the nucleus
Fig 8.5

5. Basics about an Atom Nucleus: Protons and Neutrons
Electrons (in “orbit”)
Atomic number
Isotopes
Atomic Mass Unit (u)
Atomic Weight
Mass Number
Shorthand
Fig 8.7
Subatomic particles: Protons, Neutrons, and Electrons
Atomic number
Number of protons in nucleus
Elements distinguished by atomic number
Overall Atoms are neutral
Isotopes
Same number of protons; different number of neutrons
Figure 8.7
Mass number = sum of the number of protons and neutrons
Atomic mass units (u)
1/12th of carbon-12 isotope mass
Atomic weight
Atomic mass of an element, weighted-averaged of all naturally occurring isotopes

6. The Quantum Concept Max Planck (1900) Einstein (1905) Concluded…
Introduced quantized energy
Einstein (1905)
Light made up of quantized photons
Concluded…
Higher frequency photons = more energetic photons
Equation 8.1

7. Atomic Spectra Balmer Series for Hydrogen Source of light? Fig 8.8
Continuous Spectrum
Fig 8.8
Balmer Series for Hydrogen
Like an atomic
“fingerprint”
Source of light?
Blackbody radiation
Continuous radiation distribution (Continuous Spectrum)
Depends on temperature of radiating object
Characteristic of solids, liquids and dense gases
Line spectrum from an incandescent source
Narrow lines (no light in between)
Lines of colors result from the “transitions” from energy level to energy level (n-values)
Emission at characteristic frequencies
Illustration: Balmer series of hydrogen lines
Fig 8.9: always the same colors of lines in the same position
Equation 8.2 (R is the Rydberg Constant)
Spectroscope
Device that uses a prism (or diffraction grating) to separate light into its component colors
Creates a “line” spectra because of the slit in the spectroscope
Line Spectrum

8. and their Characteristic
Some Gaseous Elements
and their Characteristic
Spectral Lines
Like Fingerprints
Think of your Fireworks displays and the colors you see

9. Example 8.3
NOTE: x 10-7 is 411 x 10-9 (or 411 nanometers, the typical units within a spectroscope)
You can then determine the “quanta” of energy of that level by:
first, determining the frequency (f) from c=λf
then, using Planck’s Law (E=hf)

10. Bohr Model
Fig 8.10
Figure FYI: Bohr won the Noble Prize for this work!
Ran with Balmer’s findings
Electrons only exist in certain allowed orbits ; more potential energy at higher orbit
Radiation is emitted or absorbed when changing orbits
Each time an electron makes a “Quantum” leap, moving from a higher-energy orbit to a lower-energy orbit, it emits a photon of a specific frequency and energy level
Bohr used the term “orbit”. A modern equivalent is “shells” or “energy levels”

11. Bohr’s “Quantum Leap” Fig 8.11
Lowest energy state = “ground state” ; AKA “normal” state (stable)
Any Higher state = “excited state” (unstable thus it doesn’t stay in that state very long)
Photon energy equals difference between energies of “states” (orbit)
Equation 8.4
Energy photon released when the electron takes a “quantum leap”
A particular color shows up
The color depends upon which levels it jumped (leaped) from and to
Hydrogen atom example (Figure 8.11)
Energy levels (in Joules) ; Don’t worry about electron volts (eV)
Line spectra

12. Example 8.4

13. Quantum mechanics (Q-M)
Bohr theory only modeled the line spectrum of Hydrogen
Main Basis of Q-M Theory: Wave nature of electrons
Main similarities between Bohr and Q-M Models:
Electrons are able to emit light (Photons)
Concept of Outer Orbitals
FYI
Louis de Broglie (1923)
Postulated matter waves
Wavelength related to momentum
Matter waves in atoms are standing waves
Wave Mechanics
Developed by Erwin Schrodinger
Treats atoms as three dimensional systems of waves
Contains successful ideas of Bohr model and much more
Describes hydrogen atom and many electron atoms
Forms our fundamental understanding of chemistry
Quantum numbers specify electronic quantum states
Visualization - wave functions and probability distributions

14. The Periodic Table of Elements
IA, IIA, etc : US Notation ; : International Notation
You must learn the “Guide” on the bottom (see red box) ; Atomic Number ; Atomic Weight
Rows = Periods
Each period begins with a single electron in a new “orbital”
Each period ends with the filling of an orbital, completing the maximum number of electrons that can occupy that main energy level
Columns = Families or Groups
The Number identifying the A families (groups) also identifies the number of electrons in the outer orbitals (exception is Helium)
Outer orbital Electrons are most responsible for the chemical properties of that element
All elements in a family or A-group have similar outer configurations
A Group (main Group) Elements: Representative Elements
Alkali Metals (IA): Very reactive metals ; especially reactive with Halogens (Video on textbook website)
Alkaline Earths (IIA): Similar to IA metals but not as reactive
Halogens (VIIA): “Salt formers” ; very reactive with IA and IIA metals
Noble Gases (VIIIA): Inert gases ; Chemically stable and inactive 8 electrons within the outer-most orbital (filled in pairs)
B-group families are called Transition Elements (Transition Metals)

15. Electron Dot Notation (AKA Lewis Dot Notation)
Representative Elements
Shorthand way to represent number of outer orbital electrons
Electrons within the Highest energy level
Relate to Group A # in the Periodic Table
Noble gases - filled electron “shells”, inert
Elements want to be Noble
The outer-most orbital electrons are responsible for chemical behavior of an atom
Fig 8.18

16. Ions Lose electron = Positive Ion Fig 8.20
Gain electron = Negative Ion
Ions: Formed by a Gain or Loss in Outer Orbital Electrons
A family of representative elements (A-Groups) all form ions with the same charge
A-Groups: 1, 2, and 3
Tend to lose electron and form positive ions (Cations)
Metals
A-Groups: 5, 6, and 7
Tend to gain electrons and form negative ions (Anions)
Nonmetals
Transition Elements (B-Group): variable charges on their Ions

17. What are their Properties?
Metals and Nonmetals
What are their Properties?
Metals (80 percent of the elements are “metals”)
Soft ; Shiny ; Malleable (hammer into sheets) ; Ductile (draw into wires)
Good conductors of electricity and heat
Have 1, 2, or 3 valence electrons
Easily lose electrons and form (+) ions (cations)
React violently with water
Nonmetals
Brittle ; non-ductile ; poor conductors
Have 4 or more valence electrons
Easily gain electrons thus can form (-) ions (anions) very easily

18. Next: Exam 3 Then Chapter 9: Chemical Bonds