Elements and Atoms - All matter (including minerals) is composed of elements (e.g. Si, O, H - simplest chemical compound that cannot be broken down to another substance by simple chemical procedures) that occur as discrete atoms (smallest form of element that maintains properties of element). Atoms contain dense nucleus of protons (positively charged) and neutrons (no charge) with electrons (negatively charged) orbiting around nucleus. Element is defined by number of protons in nucleus (atomic number).
Electron energy levels - Different electrons have discrete energy levels given by their principal quantum number, n, which varies from 1 to 7 (from low energy to high energy levels) and also is given by letters K, L, M, N, O, P, and Q. Fig. 2.7 (of handout distributed in class) shows simplified picture of different atoms (H, He, Li, and Na) with their electron energy levels (orbitals or shells); 3 circles in Na represent first, second, and third orbital.
Electron orbitals are more complicated than Fig. 2.7 indicates. Within each electron energy level, there are sublevels of electron orbitals of different shape. From low to high energy: s, p, d and f. s-orbitals are spherical and contain maximum of 2 electrons; p-orbitals are dumbbell-shaped and contain up to 6 electrons; d-orbitals are like 4-leaf clovers and contain up to 10 electrons; f-orbitals are complex and contain up to 14 electrons.
To designate energy level of electron, we combine electron level
and sub-level (e.g., 1s = lowest energy level in s-sublevel, 2s,
2p, 3s, 3p, etc.). Inner shell (1s) vs. outer shell electrons
(largest value). Electrons fill orbitals from lowest to highest
Principle). (NOTE: Above
electron energy level, rank of energy level is more complex.)
Ions - In neutral atom, number of protons (p) equals number of electrons (e). Atoms become more stable if electrons completely fill level or sublevel. To do this, atoms may either give up or capture electrons and become ions. Cations have more protons than electrons (#p > #e) and therefore have positive charges. Anions have more electrons than protons (#e > #p) and therefore have negative charges. Ionic charge (valence) = #p - #e. Usually present ions with chemical symbol followed by valence as superscript, e.g., Fe2+; which has 26 protons and 24 electrons. +1 or -1 = monovalent, +2 or -2 = divalent, +3 or -3 = trivalent. Fig. 1.7 shows common valence state of elements in minerals.
Table of Elements (link
#2 = table where each chemical element contains a link to a
page that explains its chemical properties, health effects,
environmental effects, application data, an image and also
information of the history/inventor of each element, #3) - 118
elements, 94 naturally occurring, 24 artificially produced
(transuranic) elements that form in nuclear reactors and atom
link to Periodic Table sung by Tom Lehrer and graphics)
Periodic table of elements is separated into rows (periods) and columns (groups). Period number = # of orbitals with electrons (e-'s), e.g., first period has 1s e-'s; second period has 2s and 2p (+ 1s); third period has 3s and 3p (+ lower levels); fourth and fifth periods have complex combinations of s, p, and d orbitals; sixth and seventh periods have complex combinations of s, p, d and f orbitals (lanthanides and actinides). Elements in same group (column) have similar chemical properties because they have similar # of valence e-'s. Elements on left side have a few "extra" e-'s in outer orbital and easily give them up to become cations (e.g., group 1 = alkali elements with +1 charge, H, Na, and K; group 2 = alkaline earth elements with +2 charge, Mg and Ca). Elements on right side have nearly full outer orbital and easily accept extra e-'s to become anions (e.g., group VII = halogens with -1 charge, F and Cl). At far right = inert elements (noble gases) with completely filled outer orbitals, and do not accept or give away e-'s, resulting in neutral atom (He, Ne, and Ar). Some valence states of other geologically important elements include O = 2-, Si = 4+, Al = 3+. Transition elements occupy central part of table; these elements have partially filled d- and f-orbitals and non-systematic valence states (Ti = 4+, Fe = 3+ or 2+).
Atomic Mass - Atomic number (# of protons) controls properties of element. Number of neutrons in nucleus can vary for same element, yielding different isotopes, e.g., 12C, 13C, and 14C are 3 isotopes of carbon with 6 protons and 6, 7, and 8 neutrons, respectively. Mass number = sum of p and n, so carbon mass numbers are 12, 13 and 14. Atomic weight of element = sum of masses of its naturally occurring isotopes multiplied by natural abundance of each isotope. So, mass numbers are always integers, but atomic weights are not. Atomic weight of carbon is 12.01; carbon mostly occurs as 12C isotope.