1. Big Science from the Small World of Atom - Understanding Atomic Structure with 3D Visualization By Dawen Li, Shoieb Shaik, Scott Wehby

2. Atomic Structure  Electrons: Negatively charged electrons form electron cloud surrounding nucleus  Protons: One unit positively charge.  Neutrons: No charge, along with protons forms nucleus center.  Quark: Smallest particles existed in proton and neutrons. Further discovery is still ongoing.

3. Discovery of Atomic Structure 400 B.C 1800s 1913 1904 1911 Current  Greek philosopher Democritus proposed the idea that atoms make up all substances  He named the smallest piece of matter “atomos,” meaning “not to be cut.”  However, a famous Greek philosopher, Aristotle, disputed Democritus’s theory.  Aristotle’s incorrect hypothesis was accepted for about 2000 years.  John Dalton, an English scientist, brought back the atom idea and proved atoms exist by experiment.  All elements are composed of atoms, modeled as an indivisible sphere.  Atoms of the same element are exactly alike. Atoms of different elements are different.  An English scientist J.J. Thomson proposed a so called “Plum Pudding” model: negatively charged electrons were scattered throughout a positively charged sphere, like raisins in a pudding  Another English physicist Ernest Rutherford concluded from experiments that almost all the mass of an atom, and all its positive charges, were concentrated in a central atomic nucleus surrounding by electrons.  Niels Bohr, a Danish scientist, put electrons into orbits.  He hypothesized that electrons travelled in a definite orbits around the nucleus at a specific energy level, much like planets circle the sun.  James Chadwick, a student of Rutherford, concluded that the nucleus contained positive protons and neutral neutrons.  The current electron cloud model: a cloud of electrons surrounding the dense nucleus.

4. Atomic Mass and Atomic Number Atomic Mass:  Mass of an atom is so small, even gram is extremely big measurement unit.  Nucleus contains most of the atom mass. Mass of proton = mass of neutron, both are much larger than electrons.  Mass number is the sum of # of protons and neutrons. Atomic Mass Unit = 1/12 of Carbon 12 (six protons and six neutrons) = mass of each proton or neutron Atomic Number:  Different elements have different number of protons called atomic number.  Each element can be identified by atomic number.  In an atom, # of protons = # of electrons, so atom is charge neutral. # of neutrons = mass number - # of protons

5. Isotopes and Average Atomic Mass  Average atomic mass of an element: weighted-average mass of the mixture of its isotopes.  Isotopes: atoms of the same element that have different # of neutrons

6. Valence and Core Electrons  Valence electrons are those occupying the outermost shell (the highest energy level) of an atom  Core electrons are placed in the inner (low) energy levels – close to nucleus.  Valence electrons involve chemical bonding. Therefore, these electrons are critical for chemical properties.  When an atom loses or gains an electron, it becomes either positively or negatively charged ion. Core electrons Valence electrons

7. Electron Arrangement in Shells  Electrons first fill in the lower shell (low energy level), than only move to higher one.  The maximum number of electrons in a shell = 2n 2, where n stands for the number of shell.  Each shell can further divide into sub-shells (orbitals), such as s, p, d, f orbitals.  Electrons fill orbitals in the order of 1s 2, 2s 2, 2p 6, 3s 2, 3p 6, 4s 2, 3d 10, 4p 6, 5s 2 4d 10 and so on.

8. Elements Organized by Periodic Table  In late 1800s, Russian chemist, Dmitri Mendeleev, organized all elements known at that time by atomic mass.  He found repeating patterns of chemical properties using his periodic table.  Based on periodic table, he predicted properties of missing elements, which are extremely close to what scientists discovered later on.

9. Atom and Current Periodic Table  In 1913 a young English scientist, Henry Moseley, re-arranged elements based on atomic number instead of atomic mass, which is the current periodic table.  Elements in each group (vertical column) have similar properties because they have the same number of valence electrons for chemical bonding.  The number of valence electrons can be found from I, II, … VIII, group number.  Each periodic row ends with noble gas. The outer energy level is fully filled with eight valence electrons (two electrons at row 1), leading to be “inert” for chemical reactions.  Solid metals are located in the left side of periodic table, while nonmetal gases are in right-top region.

10.  Niels Bohr, aDanish scientist, put electrons into orbits.

11. Information from Core Electrons  Moseley’s law: the difference in energy between shells changes as the atomic number varies.  When an electron transits from outer shell to the inner shell, a X-ray photon is emitted.  Energies of X-ray photons are characteristic. They are specific to the elements in the specimen, which is basis for element identification.

12. Nanotechnology from Electron Beam  Electron-beam lithography to make nano- pieces.  Electron microscopies to see nanoscale structures. Nanopillar arrays from Prof. Dawen Li’s lab at the University of Alabama

13. Discovery Using Scientific Method

14. Review Questions 14 What are different parts of an atom and draw its structure? What are the particles that make up protons and neutrons? Where is the majority of the mass of atom located? Discuss the charge from a proton, an electron, a neutron, and the charge in an atom as a whole. Which element is an atom with 14 protons in the nucleus? What is the number of valence and core electrons? Why do elements in the same group undergo similar chemical properties? How isotopes are defined? How to calculate number of neutrons? What are the properties (metal or nonmetal, solid or gas phase) of the elements located on the left side of the periodic table? Why are the noble gases in group VIII chemically stable?