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)

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): 9.11 x 10-31 kg (see front of book)

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

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

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

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

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

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

Example 8.3 NOTE: 4.11 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)

Bohr Model Fig 8.10 Figure 8.10 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”

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

Example 8.4

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

The Periodic Table of Elements IA, IIA, etc : US Notation ; 1-18: 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)

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

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

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

Next: Exam 3 Then Chapter 9: Chemical Bonds