Basic Atomic Chemistry

Early Chemistry  Early Chemists only believed in 1 element  Later Chemists believed in 4 elements: Earth Earth Air Air Fire Fire Water Water

Basic Chemistry  Atom = smallest unit of matter in universe (means “indivisible”) (means “indivisible”)  Elements are specific kinds of an atom.  Atoms are mostly empty space with a nucleus at its core and electrons at different energy levels around that nucleus.

Basic Chemistry  Atoms have electrons which are very small and are negatively charged and have a negligible mass (mass ≈ 0)  Electrons move in “orbitals” around the nucleus - at different Energy Levels.

Basic Chemistry  Atoms have a Nucleus made up of Protons & Neutrons  Protons are Positively Charged and have a mass =1 amu  Neutrons have no charge and are therefore Neutral & have a mass = 1 amu +

Summary of Subatomic Particles: Particle Name LocationChargeMass Effect/Im pact Electron Electron Cloud ≈0 Charge+ Bond (Energy) ProtonNucleus+1 1 amu Mass+ Charge NeutronNucleus No Charge 1 amu Mass+ Radi oactivity

Periodic Table Notation:  Different elements are represented on the Periodic Table using this format  The letter (symbol) is an abbrev. of Element Name  Atomic Number is ALWAYS the # of protons the atom has

Atomic Number  Determines the specific identity of an element.  Examples: Carbon = 6 Carbon = 6 Hydrogen = 1 Hydrogen = 1 Oxygen = 8 Oxygen = 8 Nitrogen = 7 Nitrogen = 7

Bohr Model  Simplifies an element and its parts. Try to make the model for Carbon, 12 C 6

“Forms” of atoms  About 100 different elements in Per. Table … but about 2000 forms of those ION : same atom w/ a varying # of electrons. Neutral atoms have EQUAL numbers of protons (+) and electrons (-) so overall charge = 0. + ions lose e’s & – ions gain e’s. + ions lose e’s & – ions gain e’s.

2. Isotopes:  Same atoms (the same atomic number) but different mass numbers b/c... different numbers of neutrons. different numbers of neutrons. Isotopes of Hydrogen 1H11H1 2H12H1 3H13H1

 Mass number is the total mass of an atom in AMU = the # of protons + neutrons.  Remember: mass number CAN change without changing the identity of the element.

Average Atomic Mass  The atomic mass on the Periodic Table is not a whole number.  When atomic mass is calculated scientists use a weighted  When atomic mass is calculated scientists use a weighted average Calculated by adding the masses of each isotope × their relative abundances

Practice Calculation  On board

Ex: Hydrogen has an average atomic mass = This means that most hydrogens have an mass of 1 but a few = 2 and even fewer = 3. Thus, the average is slightly above 1. Q: What if hydrogen’s average atomic mass = 1.95? What can we infer? MARBLE LAB

Radioactive isotopes  = unstable nucleus which decays and gives off radiation energy.  How are isotopes used in biology? 1. As tracers—organisms use elements regardless of it isotopic form, so scientists can label and follow a chemical’s path.

Suppose a runner has severe pain in his shin. The Dr. decides to check to see if the tibia has a stress fracture. The runner is given an injection w/ Technetium—99m. This radioisotope is a gamma ray producer with a half-life of 6 hours. After a several hour wait, the patient undergoes bone imaging. Any area of the body that is undergoing unusually high bone growth will show up as a stronger image on the screen. If the runner has a stress fracture, it will show up on the bone imaging scan.

2. Verifying Climate Change is in part human caused. #1 In the last 150 years we have increased the production of CO 2 via industrial waste and by deforestation (we can measure how much).

Relationship between global temperatures and CO 2 levels

How do we know it is from fossil fuels and not from the earth itself?  Industrial Revolution and beginning of CO 2 increase coincide  Carbon dating implicates humans

 CO 2 from burning fossil fuels or forests has quite a different isotopic composition from CO 2 already in the atmosphere--because plants have a preference for the lighter (faster moving) isotopes ( 12 C vs. 13 C). Thus they have lower 13 C/ 12 C ratios than that of the atmosphere.

 Since fossil fuels are derived from ancient plants, plants and fossil fuels all have roughly the same 13 C/ 12 C ratio – about 2% lower than that of the atmosphere. As CO 2 from these materials is released and mixes with the atmosphere, the average 13 C/ 12 C ratio of the atmosphere has decreased (Real Climate: Climate Science from Climate Scientists)

3. Radiometric dating (p ) Every unstable isotope decays at a constant rate called half-life = the time it takes a given amount of an isotope to decay to half of its original amount. Every unstable isotope decays at a constant rate called half-life = the time it takes a given amount of an isotope to decay to half of its original amount.

Example  Carbon dating. C-14 and C-12 are two different isotopes of carbon. The ratio of C-12 to C-14 in living tissue is relatively constant from one organism to the next. But what happens after the organism dies?  The amount of C-14 will go down compared to the amount of C-12. Why?

Practice Problems (in notes ok)  The ½ life of Carbon-14 ≈ 5730 years. How old would a fossil be if it contains 1g of C-14 compared to a living sample of equal mass that has 8 g of C-14?  What amount (% or fraction) of C-14 would you be left after 4 half-lives?  HW: p. 569 #6,7

Honors  Go to this web site: cfmhttp:// 403.cfm cfm cfm Answer these questions. 1. What does C-14 decay into? 2. If C-14 is decaying, why hasn’t it just all disappeared over time? 3. Why is C-14 not good to use when dating things less than a 10 years or more than 1 million years old?