Rutherford found by bombarding gold foil with alpha particles (42He or α) in the gold foil experiment, that most of the atom is made up of empty space. The nucleus contains most of the mass of the atom and consists of positively charged protons and neutral neutrons (Which were discovered much later by Chadwick). The masses of both these subatomic particles are 1amu. Orbiting the region around the nucleus are negatively charged electrons They have a negligible mass ( about 1/1836 of that of the proton).
An element's atomic number refers to the number of protons (nuclear charge) in the nucleus. Neutral atoms always have the same number of protons and electrons.
The sum of the mass of protons and neutrons inside the nucleus is called the mass number. One atomic mass unit is been defined as exactly 1/12 of the mass of the C-12 atom.
Gram Atomic Mass - the mass in grams of 1 mole (6.02 × 1023 particles) of an element.
For example - 12g of carbon atoms equals 1 mole of carbon atoms which equals 6.02 × 1023 carbon atoms. The atomic mass of the element always gives you the gram atomic mass of an element.
Keep in mind that the atomic masses found on your periodic tables are weighted averages. For instance carbon's mass on the periodic table is listed as 12.011 amu. This simply means that the most abundant isotope (same atomic number different number of neutrons) of carbon has 6 protons and 6 neutrons.
Bohr found that electrons in atoms can only reside at very specific energy levels. These energy levels are called principal energy levels. The energy levels are labeled 1, 2, 3, 4, 5, 6, 7... Electrons closest to the nucleus have the least amount of energy and those further from the nucleus have more energy.
Ground state - all the electrons are found in the lowest energy levels available.
Excited state - an atom absorbs energy electrons move to higher energy levels. Example - an electron usually found in the 1st energy level is now in the 2nd energy level.
Principal Quantum Number - referred to as n. It defines the principal energy level of an electron. Example - an electron in the 4th principal energy level has a principal quantum number 4.
Sublevels - Each principal energy level is broken down into sublevels. The quantum number is referred to as l (Azimuthal quantum number). Each energy level can have a total possible number of sublevels equal to its principal quantum number.
The names of the sublevels are:
The s,p,d, and f sublevels have different shapes.
Each principal energy level can have sublevels with the quantum number 0 to n-1.
In the 3rd principal energy level the possible sublevels are: 0,1,2. Therefore, in the third principal energy level you can have a s sublevel, a p sublevel, and a d sublevel.
Every principal energy level has an s sublevel.
Orbitals - a region in space where an electron can be found. Each orbital can hold only 2 electrons.
The s sublevel has 1 orbital and can hold 2 electrons
The p sublevel has 3 orbitals and can hold 6 electrons
The d sublevel has 5 orbitals and can hold 10 electrons
The f sublevel has 7 orbitals and can hold 14 electrons
The 2 represents the 2 electrons in the 1st principal energy level. The 8 refers to the 8 electrons residing in the 2nd principle energy level. The 1 refers to the 1 electron residing in the 3rd principal energy level.
Writing the electron configuration appropriately we would get 1s22s22p 63s1
Remember that on the Regents exam you are only responsible for the first method of expressing electron configuration.
Filling of orbitals
Lets take the three p orbitals for example. Each of the 3p orbitals must have one electron inside of it before a second electron may occupy the same orbital. The same goes for the s, d, and f. Each of the orbitals must have 1 electron inside the orbital before a second can occupy it. This is referred to as Hund's Rule.
The sublevels fill in this order:
Valence electrons - The number of electrons in the outermost principal energy level.
This site was last updated 10/13/02