Download Key Concepts - Chemistry Classes of Professor Alba

Key Concepts
The Atomic Theory (2.1)
The idea that all matter is composed of small, indestructible particles called atoms dates
back to the fifth century B.C.; however, at the time the atomic idea was rejected by most
Greek thinkers. At about A.D. 1800 certain observations and laws including the law of
conservation of mass, the law of constant composition, and the law of multiple
proportions led John Dalton to reformulate the atomic theory with the following
postulates: (1) each element is composed of indestructible particles called atoms; (2) all
atoms of a given element have the same mass and other properties; (3) atoms combine in
simple, whole-number ratios to form compounds; and (4) atoms of one element cannot
change into atoms of another element. In a chemical reaction, atoms change the way that
they are bound together with other atoms to form a new substance.
Although it was only 200 years ago that John Dalton proposed his atomic theory,
technology has progressed to the level where individual atoms can be imaged and moved
by techniques such as scanning tunneling microscopy (STM).
Structure of the Atom(2.2)
1. The Electron
J. J. Thomson discovered the electron in the late 1800s through experiments examining
the properties of cathode rays. He deduced that electrons were negatively charged, and
then measured their charge-to-mass ratio. Later, Robert Millikan measured the charge of
the electron, which—in conjunction with Thomson's results—led to the calculation of the
mass of an electron.
2. Radioactivity
In 1909, Ernest Rutherford probed the inner structure of the atom by working with a form
of radioactivity called alpha radiation and thereby developed the nuclear theory of the
atom. This theory states that the atom is mainly empty space, with most of its mass
concentrated in a tiny region called the nucleus and most of its volume occupied by the
relatively light electrons.
3. The Proton and the Nucleus
Early experiments showed that an atom consists of a very dense central nucleus
containing protons and neutrons, with electrons moving about the nucleus at a relatively
large distance from it. The nucleus has an overall positive charge.
Atomic Number, Mass Number, Isotopes (2.3)
Atoms are composed of three fundamental particles: the proton (1 amu, +1 charge), the
neutron (1 amu, 0 charge), and the electron (~0 amu, charge). The number of protons in
the nucleus of the atom is called the atomic number (Z) and defines the element. The sum
of the number of protons and neutrons is called the mass number (A). Atoms of an
element that have different numbers of neutrons (and therefore different mass numbers)
are called isotopes. Atoms that have lost or gained electrons become charged and are
called ions. Cations are positively charged and anions are negatively charged.
The Periodic Table (2.4)
The periodic table tabulates all known elements in order of increasing atomic number.
The periodic table is arranged so that similar elements are grouped together in columns.
Elements on the left side and in the center of the periodic table are metals and tend to lose
electrons in their chemical changes. Elements on the upper right side of the periodic table
are nonmetals and tend to gain electrons in their chemical changes. Elements located on
the boundary between these two classes are called metalloids. The atomic mass of an
element, listed directly below its symbol in the periodic table, is a weighted average of
the masses of the naturally occurring isotopes of the element. Rounded to a hole number,
it represents the Mass Number (A)
Molecules and Ions (2.5)
Chemical bonds, the forces that hold atoms together in compounds, arise from the
interactions between nuclei and electrons in atoms. In an ionic bond, one or more
electrons are transferred from one atom to another, forming a cation (positively charged)
and an anion (negatively charged). The two ions are then drawn together by the attraction
between the opposite charges. In a covalent bond, one or more electrons are shared
between two atoms. The atoms are held together by the attraction between their nuclei
and the shared electrons.
Chemical Formulas (2.6)
Compounds can be divided into two types: molecular compounds, formed between two
or more covalently bonded nonmetals; and ionic compounds, usually formed between a
metal ionically bonded to one or more nonmetals. The smallest identifiable unit of a
molecular compound is a molecule, and the smallest identifiable unit of an ionic
compound is a formula unit: the smallest electrically neutral collection of ions. Elements
can also be divided into two types: molecular elements, which occur as (mostly diatomic)
molecules; and atomic elements, which occur as individual atoms.
A compound is represented with a chemical formula, which indicates the elements
present and the number of atoms of each. An empirical formula gives only the relative
number of atoms, while a molecular formula gives the actual number present in the
molecule. Structural formulas show how the atoms are bonded together, while molecular
models show the geometry of the molecule.
Naming Inorganic Ionic and Molecular Compounds (2.7)
A flowchart for naming simple inorganic compounds is shown at the end of this section.
Use this chart to name inorganic compounds.