2.1 The Early History of Chemistry

• The processing of natural ores to produce metals for ornaments and weapons and the use of embalming fluids are just two applications of chemical phenomena that were utilized prior to 1000 B.C

• The Greeks were the first to try to explain why chemical changes occur

  • They proposed that all matter was composed of four fundamental substances: fire, air, water, and earth

  • Since the Greeks had no experiments to test their ideas, no conclusion could be reached about the divisibility of matter

• The foundations of modern chemistry were laid in the sixteenth century with the development of systematic metallurgy by Georg Bauer

• The first chemist to perform truly quantitative experiments were Robert Boyle

  • In his view, a substance was an element unless it could be broken down into two or simpler substances

  • He clung to the alchemists’ views that metals were not true elements and that away would eventually be found to change one metal onto another

• The phenomenon of combustion evoked intense interest in the seventeenth and eighteenth centuries

• Joseph Priestly was found to support vigorous combustion and was thus supposed to be low in phlogiston

2.2 Fundamental Chemical Laws

• The gases carbon dioxide, nitrogen, hydrogen, and oxygen have been discovered in the eighteenth century

• Antoine Lavoisier weighed the reactants and products of various reactions, in which he suggested that mass is neither created nor destroyed

  • Lavoisier’s verification of this law of conservation of mass was the basis for the developments in chemistry in the nineteenth century

  • He discovered that life was supported by a process that also involved oxygen and was similar in many ways to combustion

  • When the French Revolution broke out, Lavoisier had been executed on the guillotine as an enemy of the people in 1794

  • Chemistry was dominated by scientists who followed Lavoisier lead by weighed experiments to study the course of chemical reactions and determine the composition of chemical compounds

2.3 Dalton’s Atomic Theory

• Dalton’s Theory:

  • Each element is made up of tiny particles called atoms

  • The atoms of a given element are identical; the atoms of different elements are different in some fundamental ways

  • Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms

  • Chemical reactions involve reorganization of the atoms

• Dalton prepared the first table of atomic masses which were incorrect assumptions, but was an important step forward

• Avogadro’s hypothesis: At the same temperature and pressure, equal volumes of different gases contain the same number of particles

  • Unfortunately, Avogadro’s interpretations were not accepted by most chemists and many different assumptions were made about formulas and atomic masses

• Painstaking measurements were made of the masses of various elements that combined to form compounds

2.4 Early Experiments to Characterize the Atom

• The first important experiments that led to an understanding of the composition of the atom were done by J.J Thomson

  • Thomson found that when high voltage was applied to the tube, a ray he called cathode ray, was produced

  • He postulated that the ray was a stream of negatively charged particles now called electrons

  • He reasoned that since electrons could be produced from electrodes made of various types of metals, all atoms must contain electrons

  • Thomson furthered assumed that atoms also must contain some positive charge

• Nuclear atom: an atom with a dense center of positive charge with electrons moving around the nucleus at a distance that is large relative to the nuclear radius

2.5 The Modern View of Atomic Structure: An Introduction

• The simplest view of the atom is that it consists of a tiny nucleus and electrons that move about the nucleus at an average distance of about 10^-8 cm from it

• The nucleus is assumed to contain protons and neutrons

  • Two striking things about the nucleus are its small size compared with the overall size of the atom and its extremely high density

• The number of electrons possessed by a given atom affects its ability to interact with other atoms

• The atoms of different elements show different chemical behavior

• The atomic number is written as a subscript and the mass number is written as a superscript

• Isotopes show almost identical chemical properties and most elements in nature contain isotopes

2.6 Molecules and Ions

• Molecules can be represented in several different ways

  • The simplest method is the chemical formula, in which the symbols of the elements are used to indicate the types of atoms present, and subscripts are used to indicate the relative numbers of atoms

  • Examples of molecules that contain covalent bonds are hydrogen, water, oxygen, ammonia, and methane

  • Structural formulas may or may not indicate the actual shape of the molecule

• An ion is an atom or group of atoms

  • The best known ionic compound is common table salt, or sodium chloride, which forms when neutral chloride and sodium react

• Cation: positive ion

• Anion: negative ion

• Ionic bonding: the force of attraction between oppositely charged ions

• Ionic solids can consist of simple ions, as in sodium chloride, or of polyatomic ions

2.7 An Introduction to the Periodic Table

• The Periodic table shows all the known elements and gives a good deal of information about each

• Most of the elements are metals

  • Metals have characteristic physical properties such as efficient conduction of heat and electricity, malleability, and ductility

  • Metals tend to lose electrons to form positive ions

• Nonmetals lack the physical properties that characterize the metals

  • They tend to gain electrons in reactions with metals to form negative ions

  • Nonmetals often bond with each other by forming covalent bonds

  • Chlorine is a typical nonmetal

• All of the alkali metals members of Group 1A are very active elements that readily form ions with a 1+ charge when they react with nonmetals

• The alkaline earth metals of Group 2A form ions with a 2+ charge when they react with nonmetals

• The halogens, the members of group 7A, all react with metals to form salts containing ions with a 1- charge

• The noble gases of Group 8A exist under normal conditions as monatomic gases and have little chemical reactivity

• Periods: The horizontal rows of elements in the periodic table

2.8 Naming Single Compounds

• ****The solution is to adopt a system for naming compounds in which the name tells something about the composition of the compound

• Binary compounds: compounds composed of two elements

• The rules of following binary ionic compounds are: The cation is always named first and the anion second

• A monatomic cation takes its name from the name of the element

• A monatomic anion is named by taking the root of the element name and adding -ide

• In the binary ionic compounds, the metal presents forms only a single type of cation

• Oxyanions: Several series of anions contain an atom of a given element and different numbers of oxygen atoms

• Binary covalent compounds are formed between two nonmetals

  • In the naming of binary covalent compounds, the rules apply: The first element in the formula is named first, using the full element name

  • The second element is named as if it were an anion

  • Prefixes are used to denote the numbers of atoms present

  • The prefix mono- is never used for naming the first element

• Prefixes to indicate the number of atoms are used only in type 3 binary compounds

• An acid can be viewed as a molecule with one or more H+ ions attached to an anion

  • If the anion does not contain oxygen, the acid is named with the prefix hydro- and the suffix -ic

  • When the anion contains oxygen, the acidic name is formed from the root name of the anion with a suffix of -ic or -ous