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 asisotopes. 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, theatomic number (A), which is the number of
protons the atom contains, is indicated by the subscript, and the mass number (Z ), 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. Theatomic 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
1.66054
10-24 g.
Social Plugin