Atoms, Molecules and Ions - K. A. Boudreaux, Angelo State University, Lecture notes of Chemistry

Download Atoms, Molecules and Ions - K. A. Boudreaux, Angelo State University and more Lecture notes Chemistry in PDF only on Docsity! Chapter Objectives: • Learn the development of the atomic theory. • Understand the basic structure of the atom. • Understand the structure of the periodic table. • Learn how to write formulas and name ionic and binary molecular compounds. Chapter 2 Atoms, Molecules, and Ions Mr. Kevin A. Boudreaux Angelo State University CHEM 1311 General Chemistry Chemistry 2e (Flowers, Theopold, Langley, Robinson; openstax, 2nd ed, 2019) 2 The Road to the Atomic Theory You only arrive at the right answer after making all possible mistakes. The mistakes began with the Greeks. Tony Rothman, Instant Physics (1995) Nothing exists except atoms and empty space; everything else is opinion. Democritus Law of Definite Proportions • Law of Definite Proportions — All samples of a pure chemical substance, regardless of their source or how they were prepared, have the same proportions by mass of their constituent elements. (Joseph Proust, 1754-1826) 5 – Calcium carbonate, which is found in coral, seashells, marble, limestone, chalk, and San Angelo tap water, is always 40.04% by mass calcium, 12.00% carbon, and 47.96% oxygen. (We now know that this results from the fact that calcium carbonate is CaCO3.) Coral – Great Barrier Reef Marble – Lincoln Memorial Limestone Quarry Chalk – the White Cliffs of Dover Chalk – The Needles, Isle of Wight The Law of Multiple Proportions • Law of Multiple Proportions — Elements can combine in different ways to form different substances, whose mass ratios are small whole- number multiples of each other. (John Dalton, 1804) 6 Sample Mass of Mass of Compound Size Sulfur Oxygen Sulfur oxide I 2.00 g 1.00 g 1.00 g Sulfur oxide II 2.50 g 1.00 g 1.50 g mass of oxygen in sulfur oxide II per gram of sulfur mass of oxygen in sulfur oxide I per gram of sulfur = 1.50 g 1.00 g = 3 2 Dalton’s Atomic Theory • John Dalton (1766-1844) explained these observations in 1808 by proposing the atomic theory: – Each element consists of tiny indivisible (not quite) particles called atoms. – An element consists of only one type of atom, which has a mass that is characteristic of the element and is the same for all atoms of that element (not quite). – Atoms of one element differ in properties from atoms of all other elements. – Atoms combine in simple, whole-number ratios to form compounds. A given compound always has the same ratios of atoms (i.e., water is always H2O). – Atoms of one element cannot change into atoms of another element (not quite). In a chemical reaction, atoms change the way they are bound to other atoms, but the atoms themselves are unchanged. 7 The Electron • Thomson’s experiments showed that electrons were emitted by many different types of metals, so electrons must be present in all types of atoms. • Although Thomson was unable to measure the mass of the electron directly, he was able to determine the charge-to-mass ratio, e/m, -1.758820×108 C/g. – This meant that the electron was about 2000 times lighter than hydrogen, the lightest element, and atoms were thus not the smallest unit of matter. 10 The Mass of the Electron • In 1909, Robert Millikan (1868-1953) measured the charge on the electron by observing the movement of tiny ionized droplets of oil passing between two electrically charged plates. Since the e/m ratio was known from Thomson’s work, the mass of the electron could then be determined: 11 Charge of an electron: e = -1.6021810-19 C Charge to mass ratio: e/m = -1.758820×108 C/g Mass of an electron: me = 9.109389710-28 g Okay, Where’s the Positive Charge? • If there is negatively particle inside an electrically neutral atom, there must also be a positive charge. • The model for the atom that Thomson proposed (1904) was of a diffuse, positively charged lump of matter with electrons embedded in it like “raisins in a plum pudding” (a watermelon or a blueberry muffin might be a more familiar analogy). 12 The Discovery of the Nucleus • . . . instead, while most of the alpha-particles sailed through the gold foil, some were deflected at large angles, as if they had hit something massive, and some even bounced back toward the emitter. 15 It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you. — Ernest Rutherford, in E. N. da C. Andrade, Rutherford and the Nature of the Atom (1964) Plum pudding model Nuclear atom model The Nuclear Atom Model • Rutherford concluded that all of the positive charge and most of the mass (~99.9%) of the atom was concentrated in the center, called the nucleus. Most of the volume of the atom was empty space, through which the electrons were dispersed in some fashion. • The positively charged particles within the nucleus are called protons; there must be one electron for each proton for an atom to be electrically neutral. • This did not account for all of the mass of the atom, or the existence of isotopes (more later); the inventory of subatomic particles was “completed” (for the moment) by James Chadwick in 1932 [Nobel Prize, 1935], who discovered the neutron, an uncharged particle with about the same mass as the proton, which also resides in the nucleus. 16 17 The Modern View of Atomic Structure Atomic Number, Electrons • What makes elements different from each another is the number of protons in their atoms, called the atomic number (Z). All atoms of the same element contain the same number of protons. – The number of protons determines the number of electrons in a neutral atom. – Since most of the volume of the atom is taken up by the electrons, when two atoms interact with each other, it is the outermost (valence) electrons that are making contact with each other. – The number and arrangement of the electrons in an atom determines its chemical properties. Thus, the chemistry of an atom arises from its electrons. 20 Mass Number, and Isotopes • The mass number (A) is the sum of the number of protons (Z) and neutrons (N) in the nucleus of an atom: A = Z + N. • Isotopes of an element have the same # of protons, but different #’s of neutrons. – Isotopes of an element have nearly identical chemical behavior. – A nuclide is the nucleus of an element with a particular combination of protons and neutrons. – A particular nuclide can be indicated by writing the name or symbol of the atom followed by a dash and the mass number (e.g., hydrogen-1). 21 Atomic Symbols • The atomic symbol specifies information about the nuclear mass, atomic number, and charge on a particular element. Every element has a one- or two-letter symbol based on its English or Latin name. 22CO  Co !!!! Xy Never capitalized!Always capitalized! atomic number (protons) mass number (protons + neutrons) charge number of units in a molecule A Z atomic symbol Examples: Writing Element Symbols 1. Carbon-12 has how many protons? How many neutrons? How many electrons? 2. What would be the symbol for an element which has 14 protons and 15 neutrons? 3. What would be the symbol for an element which has 24 protons and 28 neutrons? 4. What would be the symbol for an element with 7 protons, 7 neutrons, and 10 electrons? 5. What would be the symbol for an element with 12 protons, 12 neutrons, and 10 electrons? 6. How many protons, neutrons, and electrons are there in 92 238U? 7. How many protons, neutrons, and electrons are in the 26 56Fe3+ ion? [TopHat] 25 Atomic Mass Units • The average atomic mass of an element is usually written underneath the element symbol on the periodic table. • The masses of atoms are measured relative to the carbon-12 isotope, which is defined as weighing exactly 12 atomic mass units (amu, or dalton, Da). – 1 amu = 1 dalton = 1.660539  10-24 g. – Protons and neutrons each weigh about 1 amu. – (Using carbon-12 as a reference allows the masses of other elements to be fairly close to whole numbers.) • The isotopic mass of a particular isotope is mass of one atom of that isotope measured in amu’s. (Hydrogen-1 = 1.007825035 amu, hydrogen-2 = 2.014101779 amu.) 26 Atomic Masses • When considering a sample of an element found in nature, we must take into account that the sample probably contains a number of different isotopes of the element. – For instance, hydrogen is mostly 1H (99.985%), but there is also a small percentage of 2H (deuterium, 0.015%). • The atomic mass (or atomic weight) of an element is the average of the masses of the naturally- occurring isotopes of that element, weighted according to the isotopes’ abundance. – This number is obtained by adding the weights of the naturally occurring isotopes multiplied by their relative abundances: 27 𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 =෍ 𝑖 (𝑓𝑟𝑎𝑐𝑡𝑖𝑜𝑛𝑎𝑙 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 × 𝑖𝑠𝑜𝑡𝑜𝑝𝑖𝑐 𝑚𝑎𝑠𝑠)𝑖 30 Are Atoms Real? Where Do Atoms Come From? 31 The Periodic Table of the Elements The Elements • All of the substances in the world are made of one or more of 118 elements, 92 of which occur naturally. • An element is a substance which cannot be chemically broken down into simpler substances. Elements are defined by the number of protons in the nucleus. • The elements are all assigned one or two letter symbols. The first letter is always capitalized, the second is never capitalized. • The names, symbols, and other information about the 118 elements are organized into a chart called the periodic table of the elements. 32 35 Periodic Properties • It has long been known that many of the elements have similar chemical properties. – Lithium, sodium, and potassium all perform the same reaction with water, 2M(s) + 2HOH(l) → 2MOH(aq) + H2(g) the only difference being the masses of the metals themselves and the vigor and speed of the reaction. Lithium slow Sodium fast Potassium warp speed 36 The Invention of the Periodic Table • In 1869 Dimitri Mendeleev published a table in which the elements that were known at the time were arranged by increasing atomic mass, and grouped into columns according to their chemical properties. The properties of the elements varied (more or less) in a periodic way in this arrangement. 37 Mendeleev’s Periodic Table Atomic Sodium Weight Chlorides Salts H 1 HCl Li 7 LiCl Be 9.4 BeCl2 B 11 BCl3 C 12 CCl4 N 14 Na3N O 16 Na2O F 19 NaF Na 23 NaCl Mg 24 MgCl2 Al 27.3 AlCl3 Si 28 SiCl4 P 31 Na3P S 32 Na2S Cl 35.5 NaCl K 39 KCl Ca 40 CaCl2 As 75 Na3As Se 78 Na2Se Br 80 NaBr • Mendeleev noticed that when he grouped the elements by their properties, there were some “holes” which he guessed corresponded to as-yet-unknown elements. • Mendeleev predicted some of the properties for two of these, eka- aluminum (?=68), and eka-silicon (?=72), which corresponded well to gallium (Ga, discovered in 1875) and germanium (Ge, 1886) Elements and the Periodic Table • The modern periodic table of the elements places the elements on a grid with 7 horizontal rows, called periods, and 18 vertical columns, called groups. – The elements are listed in order of increasing atomic number. – Two rows that are a part of periods 6 and 7 are shown beneath the table. – When they are organized in this way, there is a periodic pattern to the properties of the elements: elements in the same column (group) have similar chemical properties. – The arrangement of the elements on the periodic table is a reflection of the interior structure of the atom (more later). 40 Group Numbers • The group number can be written in a couple of different ways: – 1-18 is the IUPAC-recommended numbering system. This is more unambiguous, but less useful. – 1A-8A for tall columns, 1B-8B for short columns is the more commonly used numbering system. • In the 1A-8A columns, the column numbers represent the number of valence (outermost) electrons for the main-group elements. • The number of valence electrons are what primarily determines an atom’s chemistry. 41 42 www.angelo.edu/faculty/kboudrea/periodic.htm The Periodic Table of the Elements Other Tables Parts of the Periodic Table — Main Groups – Group 2A, the alkaline earth metals — lustrous, silvery, reactive metals. They are less reactive than the alkali metals, but are still too reactive to be found in the elemental form. – Group 7A, the halogens — colorful, corrosive nonmetals; found in nature only in compounds. – Group 8A, the noble (inert) gases — monatomic gases that are chemically stable and very unreactive. 45 Parts of the Periodic Table — Transition Metals • Transition metal groups — Groups 1B-8B (the shorter columns) — these metals exhibit a very wide range of properties, colors, reactivities, etc. • Inner transition metal groups — these elements belong between groups 3B and 4B, but are usually shown tucked underneath the main table: – Lanthanides — elements 58-71 (following the element lanthanum, La). Most of these are not commonly known, although some have industrial and research applications. (Also called the “rare earth elements.”) – Actinides — elements 90-103 (following the element actinium, Ac). Most of these elements are either highly radioactive, or are synthesized in particle accelerators. 46 Metals and Nonmetals • A jagged line on the periodic table separates the metals (left) from the nonmetals (right): • Metals are shiny, lustrous solids at room temperature (except for Hg, which is a liquid) – good conductors of electricity and heat. – malleable (can be hammered into thin sheets). – ductile (can be drawn into wire). – tend to lose electrons (oxidation) to form cations. 47 Chemical Formulas • Chemical compounds can be represented in a number of different ways: – molecular formula — subscripts show the number of each type of atom in a molecule – structural formula — lines represent the bonds that hold atoms together – ball-and-stick model — shows the 3D arrangement of atoms in a molecule – space-filling model – shows the relative sizes of the atoms 50 Molecular and Ionic Compounds • Most things that we encounter are not elements, but compounds, composed of two or more elements. – Binary compounds are composed of two elements (H2O, CH4, NH3, NaCl, CaCl2, etc.) – Diatomic compounds are composed of two atoms, which may or may not be the same (Cl2, CO, etc.) • There are two major types of chemical compounds: – molecular compounds — nonmetal + nonmetal • held together by covalent bonds that result from the sharing of pairs of electrons. – ionic compounds — metal + nonmetal • held together by ionic bonds, which result from the transfer of electrons from the metal to the nonmetal, producing ions. 51 Ions and the Periodic Table • Main group metals tend to lose electrons to form cations that have the same number of electrons as the preceding noble gas. The charge on the typical cation is the same as the group number. Group 1A: +1 Group 3A: +3 Group 2A: +2 • Main group nonmetals tend to gain electrons to form anions that have with the same number of electrons as the nearest noble gas. The charge on the typical anion is the group number minus eight. Group 5A: -3 Group 7A: -1 Group 6A: -2 • This is known as the octet rule — more later 52 55 Naming Chemical Compounds 56 Main-Group Metals • Group 1A, 2A, and 3A metals tend to form cations by losing all of their outermost (valence) electrons. • The charge on the cation is the same as the group number. • The cation is given the same name as the neutral metal atom, with the word “ion” added to the end. Group Ion Ion name Group Ion Ion name 1A H+ hydrogen ion 2A Mg2+ magnesium ion Li+ lithium ion Ca2+ calcium ion Na+ sodium ion Sr2+ strontium ion K+ potassium ion Ba2+ barium ion Cs+ cesium ion 3A Al3+ aluminum ion 57 Transition and Post-Transition Metals • Many transition and post-transition metals can form cations with more than one possible charge. – The common charges of the transition metals must be memorized. – The charges of the Group 4A and 5A metal cations are either the group number, or the group number minus two. • Common or trivial names: -ic endings go with the higher charge, -ous endings go with the lower charge. – Often, the name used is the Latin name of the element (e.g., iron = ferrum) – Fe2+ ferrous ion, Fe3+ ferric ion – Cu+ cuprous ion, Cu2+ cupric ion 60 Transition and Post-Transition Metals Ion Systematic name Common name Au3+ gold(III) ion Hg2 2+ mercury(I) ion mercurous ion Hg2+ mercury(II) ion mercuric ion Sn2+ tin(II) ion stannous ion Sn4+ tin(IV) ion stannic ion Pb2+ lead(II) ion plumbous ion Pb4+ lead(IV) ion plumbic ion Bi3+ bismuth(III) ion Bi5+ bismuth(V) ion Post- Transition Metals 61 Main-Group Nonmetals • Group 4A - 7A nonmetals form anions by gaining enough electrons to fill their valence shell (eight electrons). The charge on the anion is the group number minus eight. • The anion is named by taking the element stem and adding the ending -ide. Group Ion Ion name Group Ion Ion name 4A C4– carbide ion 6A Se2– selenide ion Si4– silicide ion Te2– telluride ion 5A N3– nitride ion 7A F– fluoride ion P3– phosphide ion Cl– chloride ion As3– arsenide ion Br– bromide ion 6A O2– oxide ion I– iodide ion S2– sulfide ion 1A H– hydride ion 62 Common Cations and Anions IA VIIIA 1 2 1 H Elements To Memorize He Hydrogen Helium 1+, 1- IIA IIIA IVA VA VIA VIIA — 3 4 6 Atomic Number 5 6 7 8 9 10 2 Li Be C Symbol B C N O F Ne Lithium Beryllium Carbon Name Boron Carbon Nitrogen Oxygen Fluorine Neon 1+ 2+ 4- Charges 3+ 4- 3- 2- 1- — 11 12 13 14 15 16 17 18 3 Na Mg Al Si P S Cl Ar Sodium Magnesium VIII Aluminum Silicon Phosphorus Sulfur Chlorine Argon 1+ 2+ IIIB IVB VB VIB VIIB 644444474444448 IB IIB 3+ 4- 3- 2- 1- — 19 20 24 25 26 27 28 29 30 33 34 35 36 4 K Ca Cr Mn Fe Co Ni Cu Zn As Se Br Kr Potassium Calcium Chromium Manganese Iron Cobalt Nickel Copper Zinc Arsenic Selenium Bromine Krypton 1+ 2+ 2+, 3+ 2+, 3+ 2+, 3+ 2+, 3+ 2+ 1+, 2+ 2+ 3- 2- 1- — 37 38 47 48 50 51 52 53 54 5 Rb Sr Ag Cd Sn Sb Te I Xe Rubidium Strontium Silver Cadmium Tin Antimony Tellurium Iodine Xenon 1+ 2+ 1+ 2+ 2+, 4+ 3+, 5+ 2- 1- — 55 56 57 79 80 82 83 86 6 Cs Ba La Au Hg Pb Bi Rn Cesium Barium Lanthanum Gold Mercury Lead Bismuth Radon 1+ 2+ 3+ 1+, 2+ 2+, 4+ 3+, 5+ — 88 89 7 Ra Ac Radium Actinium Lanthanides 92 Actinides U Uranium 65 • Some nonmetals form a series of oxyanions having different numbers of oxygens (all with the same charge). The general rule for such series is shown below. (Note that in some cases, the -ate form has three oxygens, and in some cases four oxygens. These forms must be memorized.) Polyatomic Ions — Oxyanions XOn y– stem + ate ClO3 – chlorate XOn-1 y– stem + ite ClO2 – chlorite XOn-2 y– hypo + stem + ite ClO– hypochlorite XOn+1 y– per + stem + ate ClO4 – perchlorate Xy– stem + ide (the monatomic ion) Cl– chloride 66 Polyatomic Ions — Ions Containing Hydrogens • Acid salts are ionic compounds that still contain an acidic hydrogen, such as NaHSO4. In naming these salts, specify the number of acidic hydrogens still in the salt. • The prefix bi- implies an acidic hydrogen. CO3 2– carbonate HCO3 – hydrogen carbonate, bicarbonate SO4 2– sulfate HSO4 – hydrogen sulfate, bisulfate PO4 3– phosphate HPO4 2– monohydrogen phosphate H2PO4 – dihydrogen phosphate 67 Writing Formulas of Ionic Compounds • The cation is written first, followed by the monatomic or polyatomic anion. • The subscripts in the formula must produce an electrically neutral formula unit. • The subscripts should be the smallest set of whole numbers possible. • If there is only one of a polyatomic ion in the formula, do not place parentheses around it. If there is more than one of a polyatomic ion, put the ion in parentheses, and place the subscript after the parentheses. – Remember the Prime Directive for formulas: Ca(OH)2  CaOH2! 70 Nomenclature of Ionic Compounds: Hydrates • Hydrates are ionic compounds which also contain a specific number of water molecules associated with each formula unit. The water molecules are called waters of hydration. • The formula for the ionic compound is followed by a raised dot and #H2O — e.g., MgSO4·7H2O. • They are named as ionic compounds, followed by a counting prefix and the word “hydrate” MgSO4·7H2O magnesium sulfate heptahydrate (Epsom salts) CaSO4·½H2O calcium sulfate hemihydrate BaCl2·6H2O barium chloride hexahydrate CuSO4·5H2O copper(II) sulfate pentahydrate 71 Nomenclature of Binary Molecular Compounds: Nonmetal + Nonmetal • Two nonmetals combine to form a molecular or covalent compound (i.e., one that is held together by covalent bonds, not ionic bonds). • In many cases, two elements can combine in several ways to make completely different compounds (e.g., CO and CO2). It is necessary to specify how many of each element is present within the compound. • In writing formulas, the more cation-like element (the one further to the left on the periodic table) is placed first, then the more anion-like element (the one further to the right on the periodic table). • Important exception: halogens are written before oxygen. For two elements in the same group, the one with the higher period number is placed first. 72 Nomenclature of Binary Molecular Compounds • The first element in the formula is given the element name, and the second one is named by replacing the ending of the element name with -ide. • A numerical prefix is used in front of each element name to indicate how many of that element is present. (If there is only one of the first element in the formula, the mono- prefix is dropped.) 1 mono- 4 tetra- 7 hepta- 10 deca- 2 di- 5 penta- 8 octa- 3 tri- 6 hexa- 9 nona- prefix (omit mono) name of first element stem of 2nd element + -ide prefix ^ 75 Examples: Formulas and Nomenclature 1. Write the formula for the ionic compound formed between the following pairs of species and provide a name for the compound. g. Na and sulfate __________________ h. Zn and Cl __________________ i. Mercury(I) and nitrite __________________ j. Mercury(II) and sulfite __________________ k. Chromium and S __________________ l. Silver and O __________________ 76 Examples: Formulas and Nomenclature 2. Name the following compounds. a. Ca(NO3)2 ___________________________ b. BaCO3 ___________________________ c. SO3 ___________________________ d. SnCl4 ___________________________ e. Fe2(CO3)3 ___________________________ f. AlPO4 ___________________________ g. N2O ___________________________ 77 Examples: Formulas and Nomenclature 3. Name the following compounds. a. CrO ___________________________ b. Mn2O3 ___________________________ c. NO2 ___________________________ d. NaNO2 ___________________________ e. PBr3 ___________________________ f. KHSO4 ___________________________ g. NiCl2·6H2O _________________________ 80 • Acids are compounds in which the “cation” is H+. These are often given special “acid names” derived by omitting the word “hydrogen,” adding the word “acid” at the end, and changing the compound suffix as shown below: Nomenclature of Acids Compound name Acid name stem + ate stem + ic acid stem + ite stem + ous acid stem + ide hydro + stem + ic acid HClO3 hydrogen chlorate chloric acid H2SO4 hydrogen sulfate sulfuric acid HClO2 hydrogen chlorite chlorous acid HCl hydrogen chloride hydrochloric acid oxoacids binary acids 81 Examples: Acid Nomenclature 6. Write formulas or names for the following acids. a. HCl ____________________ b. HClO2 ____________________ c. H2SO3 ____________________ d. H3PO4 ____________________ e. hydrofluoric acid ____________________ f. periodic acid ____________________ g. chloric acid ____________________ h. phosphorous acid ____________________ 82 Examples: Formulas and Nomenclature 7. Which of the following formulas and/or names is written incorrectly? a. NaSO4 b. Na2Cl c. MgNO3 d. magnesium dichloride, MgCl2 e. iron(III) phosphate, Fe3(PO4)2 f. tin(IV) sulfate, Sn(SO4)2 g. nitrogen chloride, NCl3 h. HBrO2, hypobromous acid