Atomic Structure—What We Know Today

　　Elements and Atoms　　An atom is the smallest particle of an element that retains the chemical properties of that element, and an element is defined as a substance that can’t be broken down or separated into simpler substances through a chemical reaction. Elements contain just one type of atom, and each different element contains a different type of atom. Take the element sulfur (S). A pile of sulfur (a yellow, powdery or crystallized substance) sitting on a table represents a single element—sulfur—and this pile of sulfur is made up of only one type of atom—sulfur atoms.　　Each atom, regardless of its identity, is made up of three types of subatomic particles. Protons, which are positively charged and situated at the center of the atom (also known as the atomic nucleus); neutrons, which are electrically neutral (meaning that they have no charge) and are also in the nucleus of the atom; and electrons, which are negatively charged and are situated outside the nucleus. The majority of the mass of an atom is contained in its nucleus: while electrons are about the same size as protons and neutrons, an electron has 1/837th the mass of protons or neutrons. You should also be aware that the nucleus of an atom is much, much smaller and more dense than the space occupied by an atom’s electrons—if an atom were the size of a football field, the nucleus would be the size of a flea on the 50-yard line!

　　The number of protons an atom possesses is what gives the atom its identity—all atoms of a particular element have the same number of protons in their nuclei. For example, all of the sulfur atoms in the pile of sulfur we looked at above have 16 protons in their nucleus. If they had one more proton in their nucleus, they would have a different identity—they’d be chlorine (Cl) atoms, and with one less, they’d be phosphorus (P) atoms. 　　Atoms of a given element can, however, differ in the number of neutrons they contain, and atoms of the same element that have different numbers of neutrons are known as isotopes. Most elements have at least two isotopes that occur naturally, although a few have just one. Now take a look at how atoms are usually symbolized:

　　This represents a carbon atom that has 6 protons and 6 neutrons. In this notation, the atomic number, which is the number of protons the atom contains, is indicated by the subscript, and the mass number, which is the number of the atom’s protons plus the number of its neutrons, is indicated by the superscript. Some relatively common isotopes of carbon can contain 5, 7, or 8 neutrons, so although their atomic numbers would all be 6, their mass numbers, respectively, would be 11 (6 + 5), 13 (6 + 7), and 14 (6 + 8). Isotopes can also be written as carbon-14, carbon-15, carbon-16, etc., or C-14, C-15, C-16, where the number represents the mass number of the atom. 　　The last thing you should know about the basic structure of an atom is that atoms have the same number of protons and electrons, and since protons are positively charged and electrons are negatively charged, neutral atoms have no net electrical charge.　　Example　　The atomic number of a certain element is 11, and its atomic mass number is 23. How many protons and neutrons does this atom have, and what is its chemical symbol? 　　Explanation　　If the atomic number is 11, this element is sodium and its symbol is Na. If the atomic mass number is 23, the number of neutrons is equal to 23 - 11 = 12.　　Atoms and the Periodic Table　　The day of the SAT II Chemistry exam, you will be given a periodic table to use while answering the questions. However, this periodic table will most likely be much simpler than the ones you use in class or have seen in your chemistry text. It will give you only two pieces of information for each element: the element’s atomic number and the element’s atomic weight, which is written below the element’s symbol in each box. The atomic weight of an element represents its average atomic mass based on the relative abundance of various isotopes of that element in nature. So, when we say that the atomic weight of carbon is 12.0107, we mean that the average weights of all of the isotopes of carbon that exist in nature, whether the carbon is carbon-11, -12, -13, or -14, is 12.0107.　　But what does it mean to say that the isotopes “weigh” 12.0107? 12.0107 what? Certainly not grams, or the isotopes would be a lot bigger than they are. Atomic weights have the unit amu, or atomic mass unit, and one atomic mass unit is equal to 1.6605410-24 g.