Advanced Placement Chemistry

Lesson plan #3

1 class period

Atoms, Molecules, and the Periodic Table

We see that the world around us is made of many different materials, some living, some inanimate. We see also that matter often changes, from one chemical form to another. In efforts to explain these observations, philosophers from the earliest times have speculated about the nature of the fundamental "stuff" from which the world is made.

Democritus (460 - 370 B.C.) and other early Greek philosophers thought that the material world must be made up of tiny indivisible particles that they called atomos, meaning "indivisible." Later, Plato and Aristotle formulated the notion that there can be no ultimately indivisible particles. The "atomic" view of matter faded for many centuries during which Aristotelean philosophy dominated western culture. With the beginnings of modern science in Europe during the seventeenth century; however, the idea of atoms again gained favor.

As chemists learned to measure the amounts of materials that reacted with one another to make new substances, the ground was laid for a chemical atomic theory. That theory came into being during the period 1803-1807 in the work of an English schoolteacher, John Dalton. Reasoning from a large number of observations, Dalton made the following postulates:

• Each element is composed of extremely small particles called atoms.
• All atoms of a given element are identical; the atoms of different elements are different and have different properties (including different masses).
• Atoms of an element are not changed into different types of atoms by chemical reactions; atoms are neither created nor destroyed in chemical reactions.
• Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

According to Dalton's atomic theory, atoms are the basic building block of matter. They are the smallest particles of an element that retain the chemical identity of the element.

Dalton's theory explains several simple laws of chemical combination that were known in his time.

• Law of constant composition

• In a given compound the relative number and kinds of atoms are constant.
• Based on postulate 4.

• Law of conservation of matter

• The total mass of materials present after a chemical reaction is the same as the total mass before the reaction.
• Based on postulate 3.

Dalton argued that atoms retain their identities during chemical reactions and that chemical reactions consist of the rearrangement of the atoms to give new chemical combinations.

Dalton used his theory to deduce the law of multiple proportions:

• If two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers.

• mid-1800s scientists began to study electrical discharge through partially evacuated tubes.
• A high voltage produces radiation within the tube; became known as cathode rays because it originated from the negative electrode, or cathode.
• Rays themselves cannot be seen, their movement could be detected because the rays caused certain materials, including glass, to fluoresce, or give off light.
• Scientists held differing views about the nature of the cathode rays. Experiments showed that they were deflected by electric or magnetic fields, suggesting that the rays were charged.
• One school of thought was that they might be a new form of radiation with some properties similar to those of light; that is, a new kind of wave phenomenon.
• Others believed that cathode rays might be composed of particles.
• British scientists J.J. Thomson observed many properties of the rays, including the fact that the nature of the rays is the same regardless of the identity of the cathode material.

• In a paper published in 1897 he summarized his observations and concluded that the cathode rays are not waves, but are particles with mass. Thomson's paper is generally accepted as the "discovery" of what became known as the electron.

In 1909 Robert Millikan (1868-1953) of the University of Chicago succeeded in measuring the charge of an electron by performing an experiment known as the "Millikan oil-drop experiment". [See figure 2.5, page 38] He then calculated the mass of the electron by using his value for the charge, 1.60 X 10^-19C, and Thomson's charge-to-mass ratio, 1.76 X 10⁸C/g:

Mass = 1.60 X 10^-19C/1.76 X 10⁸ C/g = 9.10 X 10^-28

Using slightly more accurate values, we obtain the presently accepted value for the mass of the electron, 9.10939 X 10^-28 g. This mass is about 2000 times smaller than that of hydrogen, the lightest atom.

• Since the time of Rutherford, physicists have learned much about the detailed composition of atomic nuclei. In the course of these discoveries, the list of particles that make up nuclei has grown long and continues to increase. As chemists, we can take a very simple view of the atom because only three subatomic particles - the proton, neutron, and electron - have a bearing on chemical behavior.
• Charge of electron: -1.602 X 10^-19 C
• Charge of proton: +1.602 X 10^-19C
• The quantity 1.602 X 10^-19 is called the electronic charge.
• For convenience, the charges of atomic and subatomic particles are usually expressed as multiples of this charge rather than in coulombs. Thus, the charge of the electron is 1-, and that of the proton is 1+. Neutrons are uncharged.

Atoms are extremely small; most have diameters between 1 X 10^-10 m and 5 X 10^-10m, or 100-500 pm. A convenient, non-SI, unit of length used to express atomic dimensions is the angstrom (� ). One angstrom = 1 X 10^-10m.

• What makes an atom of one element different from an atom of another element? The answer to this question centers on the number of protons in the nucleus of the atom: All atoms of an element have the same number of protons in the nucleus. The specific number of protons is different for different elements.
• Atoms of a given element that differ in the number of neutrons, and consequently in mass, are called isotopes. The symbol ¹²₆C or simply ¹²C (read "carbon-12") represents the carbon atom with six protons and six neutrons. The number of protons, which is called the atomic number, is shown by the subscript. The superscript is called the mass number; it is the total number of protons plus neutrons in the atom. For example, some carbon atoms contain six protons and eight neutrons and are consequently represented as ¹⁴C ("carbon-14").
• An atom of a specific isotope is called a nuclide. Thus, an atom of ¹⁴₆C is referred to as a ¹⁴₆C nuclide.

• The arrangement of element in order of increasing atomic number, with elements having similar properties placed in vertical columns, is known as the periodic table.
• Elements in a column of the periodic table are known as a group. Elements that belong to the same group often exhibit some similarities in their physical and chemical properties. For example, the "coinage metals" - copper, silver, and gold - all belong to group 1B.
• All the elements on left side and in the middle of the periodic table (except for hydrogen) are metallic elements, or metals.

• Majority of elements are metallic.
• Metals share many characteristic properties

• Luster
• High electrical conductivity
• High heat conductivity

• Solid at room temperature (except mercury).
• Separated from the nonmetallic elements by a diagonal steplike line that runs from boron (B) to astatine (At).

• At room temperature some of the nonmetals are gaseous, some are liquid, and some are solid. They generally differ from metals in appearance and other physical properties.
• Many of the elements that lie along the line that separates metals from nonmetals, such as antimony (Sb), have properties that fall between those of metals and nonmetals. These elements are often referred to as metalloids.

• A molecule is an assembly of two or more atoms tightly bound together.
• The resultant "package" of atoms behaves in many ways as a single, distinct object.

• Many elements are found in nature in molecular form; that is, two or more of the same type of atom are bound together. For example, the oxygen normally found in air consists of molecules that contain two oxygen atoms. We represent this molecular form of oxygen by the chemical formula O₂.
• Any molecule that is made up of two atoms is said to be a diatomic molecule.

• Common elements that exists as diatomic molecules at room temperature:

• Hydrogen, H₂
• Nitrogen, N₂
• Oxygen, O₂
• Fluorine, F₂
• Chlorine, Cl₂
• Bromine, Br₂
• Iodine, I₂

• Compounds that are composed of molecules are called molecular compounds and contain more than one type of atom.
• Most molecular substances that we will encounter contain only nonmetals.

• Chemical formulas that indicate the actual number and types of atoms in a molecule are called molecular formulas.
• Chemical formulas that give only the relative number of atoms of each type in a molecule are called empirical formulas.
• The molecular formula of a substance summarizes its composition but does not show how the atoms come together to form the molecule. The structural formula of a substance shows which atoms are attached to which within the molecule.

                         O

                        / \

                       H   H

                     Water, H₂O

• If electrons are removed or added to a neutral atom, a charged particle called an ion is formed.
• An ion with a positive charge is called a cation.
• A negatively charged ion is called an anion.
• In addition to simple ions such as Na⁺ and Cl^-, there are polyatomic ions such as NO₃^- and SO₄^2-. These ions consist of atoms joined ass in a molecule, but they have a net positive or negative charge.

• An ionic compound is a compound that contains positively charged ions and negatively charged ions.
• Ionic compounds are generally combinations of metals and nonmetals, as in NaCl. In contrast, molecular compounds are generally composed of nonmetals only, as in H₂O.