2.1 Atomic Theory of Matter
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1 Chapter Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter re-emerged in the early nineteenth century, championed by John Dalton.
2 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. matter is neither created nor destroyed in a chemical reaction. Antonie Lavoisier ( )
3 2.2 The Discovery of Atomic Structure Cathode Rays & Electrons Streams of negatively charged particles were found to emanate from cathode tubes, causing fluorescence. J. J. Thomson is credited with their discovery (1897). The Electron J. J. Thomson ( ) measured the charge/mass ratio of the electron to be coulombs/gram (C/g).
4 Millikan Oil-Drop Experiment Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other. Millikan Oil-Drop Experiment Robert Millikan (University of Chicago) determined the charge on the electron in Exp. value for the charge =1.602x10-19 C Electron Mass = C g C/g
5 Radioactivity Radioactivity is the spontaneous emission of radiation by an atom. It was first observed by Henri Becquerel ( ). Marie and Pierre Curie also studied it. Radioactivity Three types of radiation were discovered by Ernest Rutherford: particles particles rays
6 The Atom, circa 1900 The prevailing theory was that of the plum pudding model, put forward by Thomson. It featured a positive sphere of matter with negative electrons imbedded in it. Discovery of the Nucleus Ernest Rutherford shot particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.
7 The Nuclear Atom Since some particles were deflected at large angles, Thomson s model could not be correct. The Nuclear Atom 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.
8 Other Subatomic Particles Protons were discovered by Rutherford in Neutrons were discovered by James Chadwick in Subatomic Particles 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.
9 Symbols of Elements All atoms of the same element have the same number of protons, which is called the atomic number, Z. The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom. Isotopes Isotopes are atoms of the same element with different masses. Isotopes have different numbers of neutrons. 1 H 2 H 3 H 1 1 1
10 2.4 Atomic Weightss Atomic and molecular masses can be measured with great accuracy using a mass spectrometer. Average 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. Ex. 2.4 Naturally occurring chlorine is 75.78% 35 Cl (atomic mass amu) and 24.22% 37 Cl (atomic mass amu). Calculate the atomic weight of chlorine. Solution Atomic weight = (0.7578)( amu) + (0.2422)( amu) = amu amu = amu
11 2.5 Periodic Table The periodic table is a systematic catalog of the elements. Elements are arranged in order of atomic number. semimetal metal nonmetal Inert gas Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.
12 Groups These five groups are known by their names. Molecules Molecules are formed from atoms rather than ions. Atoms in a molecule are held together by covalent bonds.
13 Chemical Formulas 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. Molecular compounds are composed of molecules and almost always contain only nonmetals. Diatomic Molecules These seven elements occur naturally as molecules containing two atoms: Hydrogen Nitrogen Oxygen Fluorine Chlorine Bromine Iodine
14 Types of Formulas Empirical formulas give the lowest wholenumber ratio of atoms of each element in a compound. Molecular formulas give the exact number of atoms of each element in a compound. Types of Formulas Structural formulas show the order in which atoms are bonded. Perspective drawings also show the three-dimensional array of atoms in a compound.
15 Ionic and Molecular Compounds v Ions 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.
16 Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals. Writing Formulas 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 wholenumber ratio, divide them by the greatest common factor.
17 Common Cations Common Anions
18 Common ions Inorganic Nomenclature 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.
19 Naming Binary Ionic Compounds Name the following binary ionic compounds: LiBr AlCl 3 CaS Fe 2 O 3 Patterns in Oxyanion Nomenclature When there are two oxyanions involving the same element: The one with fewer oxygens ends in -ite. The one with more oxygens ends in -ate. NO 2 : nitrite; SO 2 3 : sulfite NO 3 : nitrate; SO 2 4 : sulfate
20 Patterns in Oxyanion Nomenclature Central atoms on the second row have bond to at most three oxygens; those on the third row take up to four. Charges increase as you go from right to left. Patterns in Oxyanion Nomenclature The one with the second fewest oxygens ends in -ite. ClO 2 : chlorite The one with the second most oxygens ends in -ate. ClO 3 : chlorate The one with the fewest oxygens has the prefix hypo- and ends in -ite. ClO : hypochlorite The one with the most oxygens has the prefix per- and ends in -ate. ClO 4 : perchlorate
21 Acid Nomenclature If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro-. HCl: hydrochloric acid HBr: hydrobromic acid HI: hydroiodic acid Acid Nomenclature If the anion in the acid ends in -ite, change the ending to -ous acid. HClO: hypochlorous acid HClO 2 : chlorous acid
22 Acid Nomenclature If the anion in the acid ends in -ate, change the ending to -ic acid. HClO 3 : chloric acid HClO 4 : perchloric acid Nomenclature of Binary Compounds 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).
23 Nomenclature of Binary Compounds The ending on the more electronegative element is changed to -ide. CO 2 : carbon dioxide CCl 4 : carbon tetrachloride 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. N 2 O 5 : dinitrogen pentoxide Nomenclature of Organic Compounds Organic chemistry is the study of carbon. Organic chemistry has its own system of nomenclature. The simplest hydrocarbons (compounds containing only carbon and hydrogen) are alkanes. The first part of the names just listed correspond to the number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.).
24 Nomenclature of Organic Compounds When a hydrogen in an alkane is replaced with something else (a functional group, like -OH in the compounds above), the name is derived from the name of the alkane. The ending denotes the type of compound. An alcohol ends in -ol. Homework 필수숙제 : 10, 28, 34, 36, 74, 100, 105 추가연습 12, 20, 26, 48, 56, 58, 66, 68, 72, 102
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