MHS Chemistry
Atomic Theory

Throughout history, human beings have wondered what they would see if they could just look a little closer. Looking outward from Earth, this has led to the exploration of our solar system directly, and of the rest of the entire universe through various telescopes. Looking inwards, our idea of what makes things up has changed and evolved a great deal in the past three thousand years.

One very early theory was that everything was made of four elements - Earth, Air, Fire, and Water - in various combinations. A Greek named Democritus [1][2][3][4] conjectured that every substance was made of tiny indivisible particles (called "atomon"), similar to our modern concept of a molecule.

In the past thousand years or so, alchemists, "philosophers," and scientists have isolated about 100 substances that cannot be changed into others. These are what we now know of as "elements," and their smallest building blocks are called "atoms."

In the past 100 years, scientists have discovered that atoms are not indivisble after all. Tiny amounts of electrical charge can be moved from atom to atom with particles called electrons. It was discovered that there were much heavier oppositely charged particles, now known as protons.

For a while scientists considered atoms to be tiny random mush-balls of these two pieces, like chocolate-chip cookie dough is made of chocolate chips spread evenly through dough. Then at the end of the 19th century, Ernest Rutherford discovered a third particle with no charge (named a "neutron"), and found that they could almost always be easily shot through a sample of very thin gold foil. The very few that were reflected back were evidence of the structure of the atom. Rutherford developed the "nuclear model" of the atom: an atom is composed of a very dense (positivley) charged nucleus taht contains almost all the mass, and a very thinly spread collection of electrons containing all the negative charge.

In the very early part of the twentieth century, Niels Bohr proposed that the electrons were arranged in layers or shells around the nucleus. This idea explained many interesting measurements and observations of changes of pure atoms, but no one could explain why the negatively charged electrons were not pulled into the positively charged nucleus.

In the 1920's this problem was effectively solved with the development of quantum theory.

Structure of the Atom
Atoms are divided into two reqions. In the center is the nucleus, containing just about all the mass of the atom. Surrounding this is the electron cloud, which contains only electrons. These electrons are not considered to follow an set path or orbit, so cloud is the best description.

The nucleus is composed of two particles: protons and neutrons. These two particles have roughly the same mass, but only the proton carries any charge. The neutron (true to its name) carries no electrical charge.

Atomic Number
The number of protons in a nucleus is called the atomic number (symbol Z). This number is how the periodic table is ordered, and it defines an element. In other words, every atom with exactly 2 protons in the nucleus is helium, and every helium atom has exactly 2 protons in the nucleus.

Mass Number & Isotopes
The mass of a specific atom is known as it's mass number, and is equal to the number of protons plus the number of neutrons. This is also always a whole number. An atom with a specific number of neutrons is known as an isotope. For instance, when discussing carbon in general, you are discussing all atoms with exactly six protons in the nucleus, regardless of the number of neutrons in the nucleus. When discussing carbon for the purposes of archaeological dating, scientists consider only carbon atoms with 8 neutrons in the nucleus. This isotope has a mass number of 14 (6 protons + 8 neutrons) and is refered to as "carbon-14."
formula: mass number = protons + neutrons

Atomic Weight
When a sample of an element is obtained from natural sources, it is a mix of all naturally occuring isotopes. This mix does not have a mass number because it is not one pure isotope, so it is more practical to use the atomic weight of the element. Atomic weight is the weighted average mass number of all the naturally occuring isotopes of an element. The term "weighted" means that more common isotopes have more effect on the average than less common isotopes. For example, boron occurs as two isotopes: boron-10 (about 19% of all boron atoms) and boron-11 (about 81%). The average of these two numbers is not 10.5, because there are more boron-11 atoms. The weighted average is 10.81. The atomic weights listed on the periodic table are all calculated this way, except for the very heavy nuclei (atomic number > 82) with no stable natural isotopes. These elements have their most stable isotope's mass number listed in parentheses.

Ions and Charge
Unless oftherwise stated, atoms are considered to be electrically neutral, with the same number of protons as electrons. In normal chemical reactions however, atoms may have more or less electrons than protons. These atoms are called ions in general. If there is an overall negative charge, they are called anions, and if there is an overall positive charge, they are called cations.
formula: charge = protons – electrons

The table below summarizes most of the information above. An asterisk (*) in the last column indicates information that is not required in this course. Note that "amu" stands for "atomic mass units."

Region---> Electron Cloud
Who lives there? Electrons Protons Neutrons  
Mass (actual) 9.109n28 1.673n24 1.675n24 grams*
Mass (relative) 0 1 1 amu
Charge (actual) –1.602n19 +1.602n19 0 Coulombs*
Charge (relative) –1 +1 0  
Symbol e– p+ n0  
Symbol   0e
note: in scientific notation, the letter n means "times 10 to the negative" and the letter p means "times 10 to the positive." So 0.0050 would be written as 5.0n3 and 2000 is 2p3

Atomic Number ("Z") = number of protons
Mass Number = protons + neutrons
Charge = protons - electrons
Atomic Weight = average mass of isotopes weighted by abundance

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