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Chapter 2 Atoms, Molecules and Ions. 2.1 Atomic Theory of Matter. The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton.
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2.1 Atomic Theory of Matter • The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton. • All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements. • Atoms of an element are not changed into atoms of a different element 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. • Law of Constant composition: • Also known as the law of definite proportions. • The elemental composition of a pure substance never varies. • Law of Conservation of mass: • The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place. • Dalton had no idea of atomic structure as we know it.
2.2 The Discovery of Atomic Structure • Up until the mid-1800’s, atoms were considered to be indivisible (i.e., no parts). • Experiments with cathode ray tubes by J. J Thomson (1887) led to the discovery of negatively charged particles (cathode rays) given off by all types of matter. • These particles, later named “electrons”, showed that atoms can be divided. • Thomson calculated ratio of an electron’s electrical charge to its mass as 1.76 X 108 coulombs per gram
Robert Millikan (1909) used his oil-drop experiment to determine the charge on an electron as 1.60 a 10-19 C • Electron mass was then calculated as
Radioactivity, the spontaneous emission of high-energy radiation by elements such as uranium, revealed three types of radiation: • Circa 1900: J. J. Thomson’s plum pudding model of the atom • The atom was a positive sphere of matter with electrons embedded in it.
The Nuclear Atom • Discovery of the atomic nucleus in 1911 by Ernest Rutherford. • The gold foil experiment: • Rutherford fired alpha () particles at a thin sheet of gold foil and observed the pattern of scatter of the particles. • Since some particles were deflected at large angles, Thompson’s model could not be correct.
Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom. • Most of the volume of the atom is “empty space”. • Protons were discovered in 1919 by Rutherford; neutrons in 1932 by James Chadwick.
2.3 Modern View of Atomic Structure • Protons and electrons are the only particles that have a charge. • Protons and neutrons have essentially the same mass. • The mass of an electron is so small we ignore it.
Element Symbols Elements are symbolized by one or two letters.
Atoms of the same element with different masses. Isotopes have different numbers of neutrons. 14 6 12 6 13 6 11 6 C C C C Isotopes: Average Atomic Mass: • Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. • Average mass is calculated from the isotopes of an element weighted by their relative abundances.
2.4 Atomic Weights • Atomic mass scale • Average atomic masses • The Mass Spectrometer: A direct and accurate means for determining atomic and molecular weights. (See handout)
2.5 Periodic Table: • A systematic catalog of elements. • Elements are arranged in order of atomic number. • Rows are periods, columns are groups. • Elements in the same groups have similar properties. • Nonmetals are on the right side of the periodic table (with the exception of H). • Metalloids border the stair-step line (with the exception of Al and Po). • Metals are on the left side of the chart.
Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.
Groups These five groups are known by their names.
2.6 Molecules and Molecular Compounds • Molecule: Two or more atoms bound tightly together which behaves as a single, distinct object. • Molecular compounds are composed of molecules and almost always contain only nonmetals.
A chemical formula shows the number and types of atoms in a molecular compound. • The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound. • Diatomic molecules are made up two atoms of the same element. .
Types of Formulas • Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound. • Molecular formulas give the exact number of atoms of each element in a compound.
Structural formulas show the order in which atoms are bonded. • Perspective drawings also show the three-dimensional array of atoms in a compound.
2.7 Ions and Ionic Compounds • When atoms lose or gain electrons, they become ions. • Cations are positive and are formed by elements on the left side of the periodic chart. • Anions are negative and are formed by elements on the right side of the periodic chart.
Ionic Bonds • Ionic bonds are formed between oppositely charged ions. • The ions are formed when atoms transfer one or more electrons • Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.
Writing Formulas for Ionic Compounds • Because compounds are electrically neutral, one can determine the formula of a compound this way: • The charge on the cation becomes the subscript on the anion. • The charge on the anion becomes the subscript on the cation. • If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor.
2.8 Naming Inorganic Compounds • Metal with Nonmetal (Ionic compound) • Write the name of the cation. • If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion. • If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses. • Examples:
Oxyanions • Oxyanions are polyatomic ions containing one or more oxygen atoms. • When there are two oxyanions involving the same element: • The one with fewer oxygens ends in -ite • NO2− : nitrite; SO32− : sulfite • The one with more oxygens ends in -ate • NO3− : nitrate; SO42− : sulfate • When there are four oxyanions involving the same element: • The polyatomic ion with the fewest oxygens has the prefix hypo- and ends in -ite Ex: ClO−: hypochlorite • The one with the most oxygens has the prefix per- and ends in –ate Ex: ClO4−: perchlorate • The one with the second fewest oxygens ends in -ite Ex: ClO2− : chlorite • The one with the second most oxygens ends in -ate • ClO3− : chlorate
Nomenclature of Binary Compounds • Contain two nonmetals • The less electronegative atom is usually listed first. • A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however.) • The ending on the more electronegative element is changed to -ide. • Carbon dioxide, CO2 • If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one. • Dinitrogen pentoxide, not pentaoxide Acid Nomenclature – see text or handout
