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CHAPTER TWO: ATOMS, MOLECULES, AND IONS Part one: Atomic Theory and Atomic Structure A. Atomic Theory of Matter. (Section 2.1) 1. 400 B.C. - Democritus: matter is not continuous but is composed of unimaginably tiny, discrete, indivisible particles he called “atoms.” 2. 1808 A.D. - John Dalton published first atomic theory: 3. B. a. All matter is composed of extremely small indivisible particles called atoms, that retain their identity during chemical reactions. b. All atoms of a given element have identical properties, which differ from those of other elements. c. Atoms cannot be created, destroyed, or transformed into atoms of another element. (except by nuclear reactions) d. Compounds are formed when atoms of different elements combine with each other in small whole-number ratios. e. The relative numbers and kinds of atoms are constant in a given compound. This theory explained Laws of Conservation of Matter and Laws of Definite Proportions. The Law of Definite Proportions. (applies to compounds) 1. A compound is a pure substance consisting of two or more different elements in a fixed ratio. 2. The Law of Definite Proportions states that: Different samples of any pure compound contain the same elements in the same proportion by mass. 3. Example: Water is always found to have the definite proportion 88.9% Oxygen and 11.1% Hydrogen by mass. Why is this? Chapter 2 a. Because water is composed of particles in which 1 atom of O is attached to 2 atoms of H. b. Yet atoms of Oxygen weigh 16 times as much as an atom of Hydrogen. Page 1 c. Thus the 2:1 ratio of H:O by atoms corresponds to a 2:16 ratio by mass. There are 2 parts H and 16 parts O by mass. % O = 16/18 x 100 = 88.9% O by mass % H = 2/18 x 100 = 11.1% H by mass C. Law of Multiple Proportions. 1. When two elements form more than one compound with each other, the masses of one element in these compounds for a fixed mass of the other element are in ratios of small whole numbers. 2. Example. H and O combine to form two compounds: H2O and H2O2 Amount of Oxygen per gram of H in 1st compound Amount of Oxygen per gram of H in 2nd compound 8.0 grams 16.0 grams = = whole # ratio 1/2 That this was true confirmed that compounds form by atoms combined in fixed whole number ratios. D. Structure of the Atom. 1. Chapter 2 Atoms are not the smallest particles. They are composed of three fundamental particles: a. electrons (e-) discovered by J. J. Thomson 1897 (cathode ray tube experiment) b. protons (p or p+) discovered by Goldstein in cathode ray tube c. neutrons (n or n0) discovered by Chadwick in 1932 Page 2 E. The Nuclear Atom Model. (Rutherford) (Section 2.2) 1. Early model of atomic structure. (Thomson) 2. Disproved by Rutherford’s Gold Foil Experiment (1910). a. α particles very dense, very fast. b. Expected all to pass through with minor deflections from hitting Thomson-like atoms. c. Actual result: 1.) Nearly all particles passed through the gold as if through empty space. 2.) Amazingly, a few rebounded as if hitting very dense (See Figure 2.8) Chapter 2 ⊕ charges in the foil Page 3 3. F. Conclusion: atoms are mostly empty space with a pointlike center of having virtually all the mass. ⊕ charge Atomic Masses (also called “Atomic Weights”) (AM or AW). (Section 2.4) 1. 19th century chemists systematized large body of data establishing mass combining ratios of elements. Here is an example of how it worked: suppose they had the following initial data: Mg and O combine ~3 to 2 by mass (to make oxide of Mg) H and O 1 to 8 (to make water) H and C 1 to 3 (to make marsh gas). O and C 8 to 3 (to make main oxide of carbon) 2. They then deduced a scale of relative atomic masses (traditionally called atomic weights). 3. Originally H was established as lightest element and given a value of 1 on the relative scale; later eventually chose a relative mass scale based on Carbon (easier to work with), and assigned C an Atomic Mass =12 because it was nearly 12 on the Hydrogen scale. Chapter 2 Page 4 4. Modern Atomic Masses are based on atomic mass units, assigning the Carbon-12 isotope of C to have mass of exactly 12 amu. 1 amu ≡ 1/12 (mass of 12C atom) OR: mass of one atom 12C ≡ 12 amu 5. On this scale, Atomic Mass of H is no longer 1 but 1.00794 amu. 6. Atomic Masses are given on Periodic Chart under element symbol. 7. Atomic Mass is an element’s average mass of an atom in amu (averaged over the stable isotopes of that element). 8. Periodic Table Symbol for Element: 12 Mg 24.305 Here 12 is the Atomic Number, Mg is the elemental symbol, and 24.305 is the Atomic Mass. 9. Atomic Number (Z) = number of protons in the nucleus. (Z determines which element) 10. Atomic Mass = mass of atoms of the element on a relative scale (amu’s). Future definition = number of grams of the element in one mole. G. Atomic Number. (Section 2.3) 1. Moseley (1913) - xray experiments showed: Each element differs from preceding element on the chart by having one more positive charge in its nucleus. 2. Every nucleus has integer # of protons equal to # of e- (in a neutral atom): H has one proton / He has 2 protons / Li has 3 protons / etc. 3. Chapter 2 Atomic number = Z = number of protons in the nucleus of an atom (determines that atom’s identity). Page 5 H. Neutrons. 1. Chadwick (1932) discovered neutrons by bombarding Be with α-particles. 2. Neutrons = neutral particles in the nucleus having about the same mass as protons. Therefore: nuclei contain neutrons as well as protons. I. Mass Number and Isotopes. (Section 2.4) 1. Isotopes of a given element contain the same number of protons (Z) but differ in number of neutrons in the nucleus. 2. Mass number = sum of protons and neutrons = atomic number (Z) + neutron number Hydrogen: mass number = 1 Deuterium: mass number = 2 Tritium: mass number = 3 3. Nuclide symbol: Example - two stable isotopes of Cl: 35 17 Chapter 2 Cl 37 17 Cl Z 17 17 mass number 35 37 # of neutrons 18 20 % natural abund. 75.77 24.23 mass (amu) 34.969 36.966 Page 6 4. Atomic Mass (AM) of an element is actually an average over the different stable isotopes. 35 Chlorine mass = 34.969 amu % Average mass of Cl 37 Cl Cl 36.966 amu 75.77 24.23 = 75.77 x 34.969 + 24.23 x 36.966 100 = 35.453 amu = 35.453 g/mol Average mass 35.45 amu = Atomic Mass of Cl J. Mass Spectrometry (M.S.) and Isotopic Abundance. (Section 2.4) 1. M.S. measures (charge/mass) ratio of charged particles. 2. Gaseous sample bombarded with high-energy e-, and some e- are knocked off the gas molecules creating positive ions. 3. These are focused into a beam and passed through magnetic field. 4. Field deflects ions by an angle based on: Chapter 2 a. voltage of field focusing the positive beam. b. magnetic field strength. c. masses of positive particles. d. charges of positive particles. Page 7 Figure 2.13. Mass spectrum of neon (1+ ions only). Neon consists of three isotopes, of which neon-20 is by far the most abundant (90.48%). The mass of that isotope, to five decimal places, is 19.99244 amu on the carbon-12 scale. The number by each peak corresponds to the fraction of all Ne+ ions represented by the isotope with that mass. K. The Periodic Table. (Section 2.5) 1. Based on the observation called the Periodic Law: The properties of the elements are periodic functions of their atomic numbers. 2. Examples: a. Elements Z = 2, 10, 18 have similar properties (He, Ne, Ar are chemically inert gases.) b. Elements Z = 3, 11, 19 have similar properties (Li, Na, K are chemically active metals combining with oxygen to form X2O compounds.) 3. First noted by Mendeleev and Meyer (1869). Arranged the 60 known elements in increasing order of atomic weight. (Atomic number was unknown concept then.) 4. Periodic Law works because Z also equals number of electrons in the neutral atom, and number of e determines properties. 5. Vertical columns are groups, horizontal rows are called periods. 6. Groups most strongly correlated with one another are: Chapter 2 a. Alkali metals: Group IA - Li, Na, K, Rb, Cs. (alkaline means basic) b. Alkaline earth metals: Group IIA - Be, Mg, Ca, Sr, Ba. Page 8 7. c. Halogens: Group VIIA - F, Cl, Br, I. (halogen = “salt formers”) d. Noble gases: Group O - He, Ne, Ar, Kr, Xe, Rn. (inert gases) Broader categorization of elements into metals and nonmetals: Slide 9 p.127 Figure: Trends in metallic character of A group elements. 8. Physical properties of metals and nonmetals: 9. Chemical properties of metals and nonmetals. Chapter 2 Page 9 10. Metalloids form a sort of boundary between metals and nonmetals. Some act as semiconductors = insulators at low T and conductors at high T. Part Two - Chemical Substances: Formulas and Names A. Chemical Formulas (Section 2.6) 1. Examples: O2 H2O 2. Shorthand for chemical composition of a substance. 3. Chemical formula shows the elements present in a substance and the ratio in which atoms of those elements are combined. 4. Some substances occur in molecular form, while others occur as ionic compounds. You'll need to learn to tell the difference. Chapter 2 CH4 NaCl NH3 Page 10 5. The chemical formula has a different meaning in those 2 cases. Examples: a. Water occurs as molecules = H2O. It is a molecular substance. Not only does the compound occur with a 2:1 ratio of H to O atoms, but the substance comes in particles H2O. Model of H2O. b. The salt calcium fluoride is an ionic substance with formula CaF2, meaning that, while the atoms of Ca and F are in a 1-to-2 ratio, it does not exist as molecules of CaF2, but as a crystal lattice containing 1 Ca for every 2 F atoms. Ionic lattice 6. Chemical formula for an ionic compound merely represents atom ratios. 7. Chemical formulas for molecular substances actually represent the make-up of the molecules themselves. 8. Various representations of molecular substances (See Fig. 2.18): Chapter 2 Page 11 9. Molecular models are required to represent 3-D geometric arrangements. Use HyperChem to show molecules a. ball and stick model. b. space-filling model. 10. The most stable forms of various elements in pure form: a. Noble gases - He, Ne, Ar, ... stable as individual atoms. (i.e., monatomic gases) b. Several common elements exist in most stable form as diatomic molecules - H2, N2, O2, F2, Cl2, Br2, I2, ... c. Some elements exist as polyatomic molecules: phosphorus exists as P4 (white phosphorus) sulfur as S8 ring. d. Chapter 2 Most elements in pure form exist not as molecules at all, but in large repeating arrays (crystalline solids) Page 12 B. Ionic Compounds. (Section 2.6) 1. Don’t exist as discrete molecules, but in large crystalline arrays of ions. 2. Ions are formed when electrons are: lost from metal atoms (like Na, K, Mg, Ca,...) gained by nonmetal atoms. (like Cl, Br, I, O, F,...) Na+ Sodium cation (Na missing 1 e-) Cl- Chloride anion (Cl with 1 e- added) NaCl is not a molecular entity. Exists in a lattice (See Fig. 2.21): 3. Chapter 2 Above was an example of monatomic ions. Memorize the others from Table 2.3 and 2.4. Page 13 4. 5. Chapter 2 There are several important molecular ions or polyatomic ions (See Table 2.5): a. NH4+ ammonium ion b. SO42- sulfate ion c. NO3- nitrate ion d. NO2- nitrite ion e. CO32- carbonate ion f. OH- hydroxide ion g. PO43- phosphate ion These can combine with other ions to form an ionic compound: a. NH4Cl ammonium chloride b. Na2SO4 sodium sulfate Page 14 6. Knowing the charges of various monatomic and polyatomic ions, you should be able to figure out some chemical formulas for thousands of ionic compounds (see Example 2.3 in text): a. potassium bromide K1+ b. f. SO42- Fe2(SO4)3 PO43- (NH4)3PO4 calcium carbonate marble, chalk, seashells Ca2+ not Ca2(CO3)2 but CaCO3 CO32- aluminum phosphate Al3+ C. Ag2S ammonium phosphate NH4+ e. S2- iron(III) sulfate Fe3+ d. not K1Br1 but KBr silver sulfide Ag1+ c. Br1- PO43- AlPO4 Organic Compounds (Section 2.7) 1. Chapter 2 These are molecular compounds that contain carbon combined with other elements such as hydrogen, oxygen and nitrogen. Page 15 Part Three: Naming Simple Inorganic Compounds (Section 2.8) A. Naming Binary Compounds. 1. Consist of two elements; either ionic or molecular. 2. Name more metallic element 1st and less metallic element 2nd. 3. Less metallic element named by adding “-ide” suffix to element’s stem name. 4. Example: Some binary ionic compounds containing metals that exhibit only one charged state. Formula KBr CaCl2 NaH 5. Example: Binary ionic compounds with metals that exhibit more than one stable charge; the charge of the metal is indicated by Roman numeral in parentheses. Formula Cu2O CuF2 FeS Fe2O3 6. Cation Charge +1 +2 +2 +3 Name copper(I) oxide copper(II) fluoride iron(II) sulfide iron(III) oxide Older method used “-ous” and “-ic” suffixes to indicate lower and higher ox#’s, respectively. Formula CuCl CuCl2 FeO FeBr3 Chapter 2 Name potassium bromide calcium chloride sodium hydride Cation Charge +1 +2 +2 +3 Name cuprous chloride cupric chloride ferrous oxide ferric bromide Page 16 7. Pseudobinary ionic compounds: one or more of the ions consist of more than one element but behave as simple ions. Example: hydroxide ion, OH- ; the cyanide ion, CN- ; thiocyanate ion, SCN- . Name of the anion ends in “-ide.” NH4+, is the common cation that behaves like a simple metal cation. Formula NH4I Ca(CN)2 NaOH 8. Binary molecular compounds: involve two nonmetals bonded together. Elemental proportions are indicated by using a prefix system for both elements (See Table 2.6). Formula SO2 SO3 N2O4 Chapter 2 Name ammonium iodide calcium cyanide sodium hydroxide Name sulfur dioxide sulfur trioxide dinitrogen tetroxide Formula Cl2O7 CS2 As4O6 Name dichlorine heptoxide carbon disulfide tetraarsenic hexoxide Page 17 9. Binary acids dissolved in water. When pure, named as typical binary compounds. Their aqueous solutions are named with the prefix “hydro-” and the suffix “-ic” followed by the word “acid.” Formula HCl HF H2S HCN B. Name of Compound hydrogen chloride hydrogen fluoride hydrogen sulfide hydrogen cyanide Name of Aqueous Solution hydrochloric acid, HCl(aq) hydrofluoric acid, HF(aq) hydrosulfuric acid, H2S(aq) hydrocyanic acid, HCN(aq) Naming Ternary Acids and Their Salts. 1. Ternary compound consists of three elements. 2. Ternary acids (oxoacids) are compounds of hydrogen, oxygen, and (usually) a nonmetal. (e.g. H2SO4) 3. Nonmetals that exhibit more than one stable charge form more than one ternary acid, differing in number of oxygen atoms. 4. Most common ternary acid is given the “-ic” name. See following. Examples: Nitric acid Sulfuric acid Chapter 2 Page 18 Phosphoric acid Chloric, Bromic, and Iodic acids Carbonic acid 5. Acids containing one fewer oxygen atom per central atom are named “-ous.” Formula H2SO3 HNO2 6. Chapter 2 Name hypochlorous acid hypophosphorous acid Name perchloric acid perbromic acid periodic acid The oxoacids of chlorine follow as example: Formula HClO HClO2 HClO3 HClO4 9. Name chlorous acid Acids containing one more oxygen atom per central nonmetal atom than the normal “-ic acid” are named “per” “ic” acids. Formula HClO4 HBrO4 HIO4 8. Formula HClO2 Acids that have fewer O atom than the “-ous” acids are named using the prefix “hypo-” and the suffix “-ous.” Formula HClO H3PO2 7. Name sulfurous acid nitrous acid Ox. No. of Cl +1 +3 +5 +7 Name hypochlorous acid chlorous acid chloric acid perchloric acid Ternary salts: for example, KClO3 - potassium chlorate. a. Anion derived from “-ic acid” is “-ate.” b. Anion derived from “-ous acid” is “-ite.” Page 19 c. The “-per-” and “hypo-” prefixes are retained. Examples: HClO3 chloric acid KClO3 potassium chlorate ClO3chlorate ion “ic acid” → “ate” HClO2 chlorous acid NaClO2 sodium chlorite ClO2chlorite ion “ous acid” → “ite” HClO hypochlorous acid HClO4 perchloric acid NH4ClO ammonium hypochlorite KClO4 potassium perchlorate ClOhypochlorite ion ClO4perchlorate ion 10. Summary Chart: naming ternary acids and their anions. The stem (XXX) represents the stem of the name, e.g., “nitr,” “sulfur,” or “chlor.” 11. Examples: Formula (NH4)2SO4 KNO3 Ca(NO2)2 LiClO4 FePO4 NaClO Chapter 2 Name ammonium sulfate potassium nitrate calcium nitrite lithium perchlorate iron(III) phosphate sodium hypochlorite Page 20 12. Ternary acids salts in which one or more acidic hydrogen atoms remain: named with the word “hydrogen” or “dihydrogen” inserted after the cation. Formula NaHSO4 NaHSO3 KH2PO4 K2HPO4 NaHCO3 Name sodium hydrogen sulfate sodium hydrogen sulfite potassium dihydrogen phosphate potassium hydrogen phosphate sodium hydrogen carbonate (sodium bicarbonate) Part Four: Chemical Reactions: Equations A. Writing Chemical Equations. (Section 2.9) 1. 2. Shorthand for a chemical reaction, showing: a. substances reacting. (reactants) b. substances formed. (products) c. relative amounts involved. (balancing coefficients) Example: combustion of natural gas, methane. (CH4) CH4 + 2 O2 → CO2 + 2 H2O a. reactants - CH4, O2 b. products - CO2, H2O c. Reads: 1 molecule CH4 reacts with 2 molecules O2 producing 1 molecule CO2 and 2 molecules H2O. d. Note - both sides of equation contain: 1 atom C 4 atoms H 4 atoms O Therefore, equation is balanced. Chapter 2 Page 21 e. Note - smallest whole-number coefficients are typically used to balance: 1:2:1:2 f. Note - mole interpretation is also valid. 1 mole CH4 reacts with 2 moles O2 producing 1 mole CO2 and 2 moles H2O. g. Species formulas in the equation must describe them as they exist. Example: this is improper. CH4 + 4 O → CO2 + 2 H2O This balances, but oxygen not shown as diatomic molecule. B. Balancing Chemical Equations. (Section 2.10) 1. Problem: Write down and balance the chemical equation for the combustion of propane, C3H8, in the presence of abundant Oxygen. __ C3H8 + __ O2 → __ CO2 + __ H2O a. First balance the elements that appear in only one species on both sides of the equation (C and H) b. Then complete the balance of O C3H8 + 5 O2 → 3 CO2 + 4 H2O 2. All combustion reactions can be done this way. a. Try combustion of ethanol, C2H6O. __ C2H6O + __ O2 → __ CO2 + __ H2O b. Incomplete combustion of methane: (incomplete combustion refers to production of CO rather than CO2) __ CH4 + __ O2 → __ CO + __ H2O Chapter 2 Page 22