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Chapter 2 Atoms, Molecules, and Ions ▶ A SECTION THROUGH A GEODE. A geode is a mass of mineral matter (often containing quartz) that accumulates slowly within the shell of a roughly spherical, hollow rock. Eventually, perfectly formed crystals may develop at a geode’s center. The colors of a geode depend upon its composition. Here, agate crystallized out as the geode formed. 1 What’s Ahead Atoms, Molecules, and Ions 2.1 THE ATOMIC THEORY OF MATTER We begin with a brief history of the notion of atoms—the smallest pieces of matter. 2.2 THE DISCOVERY OF ATOMIC STRUCTURE We then look at some key experiments that led to the discovery of electrons and to the nuclear model of the atom. 2.3 THE MODERN VIEW OF ATOMIC STRUCTURE We explore the modern theory of atomic structure, including the ideas of atomic numbers, mass numbers, and isotopes. 2.4 ATOMIC WEIGHTS We introduce the concept of atomic weights and how they relate to the masses of individual atoms. 2.5 THE PERIODIC TABLE We examine the organization of the periodic table, in which elements are put in order of increasing atomic number and grouped by chemical similarity. 2 What’s Ahead Atoms, Molecules, and Ions 2.6 MOLECULES AND MOLECULAR COMPOUNDS We discuss the assemblies of atoms called molecules and how their compositions are represented by empirical and molecular formulas. 2.7 IONS AND IONIC COMPOUNDS We learn that atoms can gain or lose electrons to form ions. We also look at how to use the periodic table to predict the charges on ions and the empirical formulas of ionic compounds. 2.8 NAMING INORGANIC COMPOUNDS We consider the systematic way in which substances are named, called nomenclature, and how this nomenclature is applied to inorganic compounds. 2.9 SOME SIMPLE ORGANIC COMPOUNDS We introduce organic chemistry, the chemistry of the element carbon. 3 Atoms, Molecules, and Ions How Do We Explain? • What makes diamonds transparent and hard, while table salt is brittle and dissolves in water? • Why does paper burn, and why does water quench fires? • Where does the beautiful colors of flowers come from? • The structure and behavior of atoms are key to understanding the properties of matter. • In this chapter we examine the basic structure of atoms and discuss the formation of molecules and ions, thereby providing a foundation for exploring chemistry more deeply in later chapters. 4 2.1 THE ATOMIC THEORY OF MATTER History • Democritus and Greek philosophers (BC 400). – The material world must be made up of tiny indivisible particles. – Atomos: indivisible or uncuttable. • Plato and Aristotle. – There can be no ultimately indivisible particles. – The “atomic” view of matter faded for many centuries. • The notion of atoms reemerged in Europe during the 17th century. – Chemists learned to measure the amounts of elements that reacted with one another to form new substances. – The ground was laid for an atomic theory that linked the idea of elements with the idea of atoms. 5 2.1 THE ATOMIC THEORY OF MATTER Dalton’s Atomic Theory • The theory came from the work of John Dalton during the period from 1803 to 1807. • Atoms are the fundamental building blocks of matter. • The theory was based on the four postulates given in the figure. 6 2.1 THE ATOMIC THEORY OF MATTER Dalton’s Postulates 1) Each element is composed of extremely small particles called atoms. 7 2.1 THE ATOMIC THEORY OF MATTER Dalton’s Postulates 2) 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. 8 2.1 THE ATOMIC THEORY OF MATTER Dalton’s Postulates 3) 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. 9 2.1 THE ATOMIC THEORY OF MATTER Dalton’s Postulates 4) Atoms of more than one element combine to form compounds; a given compound always has the same relative number and kind of atoms. 10 2.1 THE ATOMIC THEORY OF MATTER Dalton’s Postulates • Dalton’s theory explains several simple laws of chemical combination. • The law of constant composition. – In a given compound, the relative numbers and kinds of atoms are constant. • The law of conservation of mass (matter). – The total mass of materials present after a chemical reaction is the same as the total mass present before the reaction. • 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 (H2O and H2O2). 11 2.2 THE DISCOVERY OF ATOMIC STRUCTURE Atomic Structure • Dalton had no direct evidence for the existence of atoms. • Scientists have developed methods for more detailed probing of the nature of matter. • Many discoveries led to the fact that the atom itself was made up of smaller (subatomic) particles. • We can measure the properties of individual atoms and even provide images of them. Figure 2.2 An image of the surface of silicon. The image was obtained by a technique called scanning tunneling microscopy. The color was added to the image by computer to help distinguish its features. Each red sphere is a silicon atom. 12 2.2 THE DISCOVERY OF ATOMIC STRUCTURE The Electron (Cathode Rays) • Streams of negatively charged particles were found to emanate from cathode tubes, causing fluorescence. • J. J. Thomson is credited with their discovery (1897). 13 2.2 THE DISCOVERY OF ATOMIC STRUCTURE The Electron Thomson measured the charge/mass ratio of the electron to be 1.76 × 108 coulombs/gram (C/g). 14 2.2 THE DISCOVERY OF ATOMIC STRUCTURE Millikan Oil-Drop Experiment • In 1909, Robert Millikan (1868–1953) succeeded in measuring the charge of an electron (1.602 × 10–19 C) by performing the oil-drop experiment. Figure 2.5 Millikan’s oil-drop experiment used to measure the charge of the electron. Small drops of oil were allowed to fall between electrically charged plates. The drops picked up extra electrons as a result of irradiation by X-rays and so became negatively charged. Millikan measured how varying the voltage between the plates affected the rate of fall. From these data he calculated the negative charge on the drops. Because the charge on any drop was always some integral multiple of 1.602 × 10–19 C, Millikan deduced this value to be the charge of a single electron. 15 2.2 THE DISCOVERY OF ATOMIC STRUCTURE Millikan Oil-Drop Experiment • He then calculated the mass of the electron by using his experimental value for the charge (1.602 × 10–19 C) and Thomson’s charge-to-mass ratio (1.76 × 108 C/g): 16 2.2 THE DISCOVERY OF ATOMIC STRUCTURE Radioactivity • Radioactivity is the spontaneous emission of highenergy radiation by an atom. • It was first observed by Henri Becquerel. • Marie and Pierre Curie also studied it. • Its discovery showed that the atom had more subatomic particles and energy associated with it. Figure 2.6 Marie Sklodowska Curie (1867–1934). When Marie Curie presented her doctoral thesis, it was described as the greatest single contribution of any doctoral thesis in the history of science. In 1903 Henri Becquerel, Maire Curie, and her husband, Pierre, were jointly awarded the Nobel Prize in Physics for their pioneering work on radioactivity (a term she introduced). In 1911 Marie Curie won a second Nobel Prize, this time in chemistry for her discovery of the elements polonium and radium. 17 2.2 THE DISCOVERY OF ATOMIC STRUCTURE Radioactivity • Three types of radiation were discovered by Ernest Rutherford: – α particles: 2+ charge. • A mass about 7400 times that of an electron. – β particles: 1– charge. • A mass of an electron. – γ rays: No charge, high energy radiation. • Not consist of particles. Figure 2.7 Behavior of alpha (α), beta (β), and gamma (γ) rays in an electric field. 18 2.2 THE DISCOVERY OF ATOMIC STRUCTURE 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. Figure 2.8 J. J. Thomson’s plum-pudding model of the atom. Ernest Rutherford proved this model wrong. 19 2.2 THE DISCOVERY OF ATOMIC STRUCTURE Discovery of the Nucleus • Ernest Rutherford shot α particles at a thin sheet of gold foil and observed the pattern of scatter of the particles. Figure 2.9 Rutherford’s α-scattering experiment. 20 2.2 THE DISCOVERY OF ATOMIC STRUCTURE The Nuclear Atom • Since some particles were deflected at large angles, Thompson’s model could not be correct. Figure 2.10 The structure of the atom. A cloud of rapidly moving electrons occupies most of the volume of the atom. The nucleus occupies a tiny region at the center of the atom and is composed of the protons and neutrons. The nucleus contains virtually all the mass of the 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. • Protons were discovered by Rutherford in 1919. • Neutrons were discovered by James Chadwick in 1932. 21 2.3 THE MODERN VIEW OF ATOMIC STRUCTURE Subatomic Particles • Protons (+1) and electrons (–1) are the only particles that have a charge; neutrons are neutral. – Every atom has an equal number of protons and electrons, so atoms have no net charge. • Protons and neutrons have essentially the same mass. • The mass of an electron is so small when compared to that of a proton or a neutron. We ignore it. • Protons and neutrons are found in the nucleus; electrons travel around the nucleus. 22 2.3 THE MODERN VIEW OF ATOMIC STRUCTURE Atomic Mass • Atoms are very tiny, so a mass scale on the atomic level is used to express the mass of an atom or molecule. • An atomic mass unit (amu) is the base unit. • 1 amu = 1.66054 × 10–24 g. • The SI symbol used for the atomic mass unit is u. 23 24 2.3 THE MODERN VIEW OF ATOMIC STRUCTURE Symbols of Elements • What makes an atom of one element different from an atom of another element: A characteristic number of protons of the atoms. • Atomic number: The number of protons in the nucleus. It is written as a subscript BEFORE the symbol. • Mass number: The total number of protons and neutrons in the atom. It is written as a superscript BEFORE the symbol. • Elements are represented by a one or two letter symbol. This is the symbol for carbon. 25 2.3 THE MODERN VIEW OF ATOMIC STRUCTURE Isotopes • Isotopes: Atoms of the same element with different masses. • Isotopes have different numbers of neutrons. • The symbol (read “carbon twelve,” carbon-12) can be represented simply as . 26 27 Atomic Mass 2.4 ATOMIC WEIGHTS • Atomic and molecular masses can be measured with great accuracy using a mass spectrometer. Figure 2.11 A mass spectrometer. Cl atoms are introduced at A and are ionized to form Cl+ ions, which are then directed through a magnetic field. The paths of the ions of the two Cl isotopes diverge as they pass through the field. Figure 2.12 Mass spectrum of atomic chlorine. The fractional abundances of the isotopes 35Cl and 37Cl are indicated by the relative signal intensities of the beams reaching the detector of the mass spectrometer. 28 Atomic Mass 2.4 ATOMIC WEIGHTS • Atomic mass unit (amu). – Defined by assigning a mass of exactly 12 amu to an atom of 12C. • • 1H: 1.6735 × 10–24 g = 1.0078 amu. 16O: 2.6560 × 10–23 g = 15.9949 amu. – 1 amu = 1.66054 × 10–24 g and 1 g = 6.02214 × 1023 amu. • The average atomic mass of an element = The element’s atomic weight. – Most elements occur in nature as mixtures of isotopes. – Calculated from the isotopes of an element weighted by their relative abundances: – C: 98.93% 12C and 1.07% 13C. Atomic weight = (0.9893)(12 amu) + (0.0107)(13.00335 amu) = 12.01 amu. 29 30 Periodic Table 2.5 THE PERIODIC TABLE • The periodic table is the most significant tool that chemists use for organizing and remembering chemical facts. Figure 2.14 Periodic table of the elements. 31 Periodic Table 2.5 THE PERIODIC TABLE • The periodic table is a systematic catalog of the elements. • Elements are arranged in order of atomic number. 32 Periodic Table 2.5 THE PERIODIC TABLE • The rows on the periodic chart are periods. • Columns are groups. • Elements in the same group have similar chemical properties. – “Coinage metals” (Cu, Ag, and Au) belong to group 11 (or 1B): Less reactive than most metals. 33 Periodicity 2.5 THE PERIODIC TABLE • Many elements show very strong similarities to one another. • When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities. Figure 2.13 Arranging elements by atomic number reveals a periodic pattern of properties. This pattern is the basis of the periodic table. 34 Groups 2.5 THE PERIODIC TABLE • These five groups are known by their names. 35 2.5 THE PERIODIC TABLE Metals, Nonmetals, and Metalloids • 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, Po, and At). • Metals are on the left side of the chart. 36 2.5 THE PERIODIC TABLE Metals, Nonmetals, and Metalloids • Nonmetals generally differ from the metals in appearance and in other physical properties. • A metalloid is a chemical element with properties that are in-between or a mixture of those of metals and nonmetals. Figure 2.15 Examples of metals and nonmetals. 37 38 2.6 MOLECULES AND MOLECULAR COMPOUNDS Molecules • Only the noble-gas elements are normally found in nature as isolated atoms. • A molecule is an assembly of two or more atoms tightly bound together. • Molecules behave in many ways as a single, distinct object. Figure 2.18 Molecular models. Notice how the chemical formulas of these simple molecules correspond to their compositions. • The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule. 39 2.6 MOLECULES AND MOLECULAR COMPOUNDS Molecules and Chemical Formulas • Many elements are found in nature in molecular form. – Two different molecular forms of oxygen. • O2: A diatomic molecule, essential for life, odorless. • O3: A triatomic molecule, toxic, pungent smell. • Molecular compounds are composed of molecules and almost always contain only nonmetals. • These seven elements occur naturally as molecules containing two atoms (diatomic molecules): Hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine. 40 2.6 MOLECULES AND MOLECULAR COMPOUNDS Molecular and Empirical Formulas • Molecular formulas give the exact number of atoms of each element in a compound. – – – – H 2O H 2 O2 C 2H 4 C6H12O 6 • Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound. – – – – H 2O HO CH2 CH2O • Why do we need empirical formulas? – Certain common methods of analyzing substances (elemental analysis) lead to the empirical formula only. 41 42 2.6 MOLECULES AND MOLECULAR COMPOUNDS Picturing Molecules • A structural formulas show the order in which atoms are attached. They do NOT depict the threedimensional shape of molecules. • A perspective drawings show the three-dimensional order of the atoms in a compound. These are also demonstrated using models. • A Ball-and-stick models shows the accurate angles between bonds. • A space-filling model shows the relative sizes of the atoms. 43 2.7 IONS AND IONIC COMPOUNDS Cations and Anions • 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 (metal atoms). – Anions are negative and are formed by elements on the right side of the periodic chart (nonmetal atoms). 44 45 2.7 IONS AND IONIC COMPOUNDS Predicting Ionic Charges • Many atoms gain or lose electrons to end up with the same number of electrons as the noble gas closest to them in the periodic table. – Na Na+ (the same number of electrons as in Ne). – Cl Cl– (the same number of electrons as in Ar). • The periodic table is very useful for remembering ionic charges. – – – – – The group 1A elements (alkali metals) form 1+ ions. The group 2A elements (alkaline earths) form 2+ ions. The group 7A elements (halogens) form 1– ions. The group 6A elements form 2– ions. Many of the other groups do not lend themselves to such simple rules. 46 47 2.7 IONS AND IONIC COMPOUNDS Ionic Compounds • A compound made up of cations and anions. • Ionic compounds (such as NaCl) are generally formed between metals and nonmetals. – Cf) Molecular compounds (such as H2O) are generally composed of nonmetals only. • Electrons are transferred from the metal to the nonmetal. The oppositely charged ions attract each other. Only empirical formulas are written. 48 49 2.7 IONS AND IONIC COMPOUNDS Writing Empirical 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 whole-number ratio, divide them by the greatest common factor. • The ionic compound formed from Mg2+ and N3–: 50 51 2.7 IONS AND IONIC COMPOUNDS ELEMENTS REQUIRED BY LIVING ORGANISMS Figure 2.20 Elements essential to life. • More than 97% of the mass of most organisms: O, C, H, N, P, and S. • At least 70% of the mass of most cells: H2O. • C is the most prevalent element by mass. 52 2.8 NAMING INORGANIC COMPOUNDS Names and Formulas of Ionic Compounds • Cations. – Cations formed from metal atoms have the same name as the metal: – If a metal can form cations with different charges, the positive charge is indicated by a Roman numeral in parentheses following the name of the metal: • Ions of the same element that have different charges exhibit different properties, such as different colors. Figure 2.21 Different ions of the same element have different properties. Both substances shown are compounds of iron. The substance on the left is Fe3O4, which contains Fe2+ and Fe3+ ions. The substance on the right is Fe2O3, which contains Fe3+ ions. 53 2.8 NAMING INORGANIC COMPOUNDS Names and Formulas of Ionic Compounds • Metals that form only one cation: Metals of group 1A and 2A, Al3+, Ag+, Zn2+. • An older method uses the endings -ous and -ic added to the root of the element’s Latin name: – Cations formed from nonmetal atoms have names that end in –ium: 54 2.8 NAMING INORGANIC COMPOUNDS Names and Formulas of Ionic Compounds 55 2.8 NAMING INORGANIC COMPOUNDS Names and Formulas of Ionic Compounds • Anions. – The names of monatomic anions are formed by replacing the ending of the name of the element with -ide: – A few polyatomic anions also have names ending in -ide: – Polyatomic anions containing oxygen have names ending in either ate or -ite and are called oxyanions. • The -ate is used for the most common or representative oxyanion of an element. • The -ite is used for an oxyanion that has the same charge but one O atom fewer: 56 2.8 NAMING INORGANIC COMPOUNDS Names and Formulas of Ionic Compounds • Prefixes are used when the series of oxyanions of an element extends to four members, as with the halogens. The prefix per- indicates one more O atom than the oxyanion ending in -ate; hypo- indicates one O atom fewer than the oxyanion ending in -ite: Figure 2.22 Procedure for naming anions. The first part of the element’s name, such as “chlor” for chlorine or “sulf” for sulfur, goes in the blank. 57 2.8 NAMING INORGANIC COMPOUNDS Names and Formulas of Ionic Compounds Figure 2.23 Common oxyanions. The composition and charges of common oxyanions are related to their location in the periodic table. – Anions derived by adding H+ to an oxyanion are named by adding as a prefix the word hydrogen or dihydrogen, as appropriate: • An older method uses the prefix bi-: The HCO3– ion is commonly called the bicarbonate ion, and HSO4– is sometimes called the bisulfate ion. 58 59 2.8 NAMING INORGANIC COMPOUNDS Names and Formulas of Ionic Compounds • Ionic compounds. – Names of ionic compounds consist of the cation name followed by the anion name: 60 61 2.8 NAMING INORGANIC COMPOUNDS Names and Formulas of Acids • If the anion in the acid ends in -ide, change the ending to ic acid and add the prefix hydro-. • If the anion in the acid ends in -ate, change the ending to ic acid. If the anion in the acid ends in -ite, change the ending to -ous acid. 62 63 2.8 NAMING INORGANIC COMPOUNDS Nomenclature of Binary Molecular Compounds • The procedures used for naming binary (two-element) molecular compounds. 1. The name of the element farther to the left in the periodic table (less electronegative atom, closest to the metals) is usually written first. • An exception occurs when the compound contains oxygen and chlorine, bromine, or iodine (any halogen except fluorine), in which case oxygen is written last. 2. If both elements are in the same group, the one closer to the bottom of the tables is named first. 3. The name of the second element is given an -ide ending. • CO2: carbon dioxide, CCl4: carbon tetrachloride. 64 2.8 NAMING INORGANIC COMPOUNDS Nomenclature of Binary Molecular Compounds 4. Greek prefixes are used to indicate the number of atoms of each element. The prefix mono- is never used with the first element. When the prefix ends in a or o and the name of the second element begins with a vowel, the a or o of the prefix is often dropped. 65 66 2.9 SOME SIMPLE ORGANIC COMPOUNDS Nomenclature of Organic Compounds • Organic chemistry. – The study of carbon. – Organic chemistry has its own system of nomenclature. – Organic compounds: Compounds that contain carbon and hydrogen. • Alkanes (the simplest hydrocarbons). – Compounds containing only carbon and hydrogen. – The first part of the names just listed correspond to the number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.). 67 2.9 SOME SIMPLE ORGANIC COMPOUNDS Some derivatives of Alkanes • When a hydrogen in an alkane is replaced with something else (a functional group, like -OH), the name is derived from the name of the alkane. • The ending denotes the type of compound. – An alcohol is obtained by replacing an H atom with a -OH group. – An alcohol ends in -ol. • Compounds with the same molecular formula but different arrangements of atoms are called isomers. 68 69 Chapter 2. Homework Exercises 2.3 2.13 2.17 2.22 2.31 2.39 2.41 2.50 2.61 2.70 2.79 2.85 2.111 70