notes_Week_02.pdf

What are the two important parts of a measurement? 1 1 0% 0%0%0%1. The number and the unit 2. The number and the exponent 3. The number and the decimal 4. The decimal and exponent Which of the following is a frequently used unit for the density of liquids and solids? 2 1 2 3 4 5 0% 0% 0%0%0% 1. g/L 2. kg/L 3. g/mL 4. cm3/mL 5. g/m3 Convert 0.000459 to scientific notation? 3 1 2 3 4 5 0% 0% 0%0%0% 1. 4.59 x 104 2. 0.459 x 103 3. 4.59 x 10?3 4. 4.59 x 10?4 5. 4.59 x 10?5 Convert 3.23 x 103 to decimal notation 4 1. 2. 3 4. 5 0% 0% 0%0%0% 1. 3.23 2. 32.3 3. 323 4. 0.00323 5. 3230 Lecture 2 Chapter 2 Atoms, Molecules, and Ions 5 Fundamental Chemical Laws ? Law of conservation of mass ?Matter can either created nor destroyed in the chemical reactions ? Law of definite proportions ?A given compound always contains exactly the same proportions of elements by mass ? Law of multiple proportions ?the same elements can be combined to form different compounds by combining the elements in different proportions 6 Atoms and Molecules ? Atoms are the tiny particles that make up all matter. ? In most substances, the atoms are joined together in units called molecules. ? Molecules ? is a bonded collection of two or more atoms of the same or different elements 7 Radioactivity is a spontaneous emission of radiation 8 Alpha Radiation ?composed of 2 protons and 2 neutrons ?+2 charge ?mass of 4 amu ?helium-4 nucleus (He-4) ?creates element with atomic number 2 lower 9 Beta Radiation ?composed of a high energy electron which was ejected from the nucleus ??neutron? converted to ?proton? ?very little mass ?-1 charge ?creates element with atomic number 1 higher 10 Gamma Radiation ?nucleus has energy levels ?energy released from nucleus as the nucleus changes from higher to lower energy levels ?no mass ?no charge 11 Structure of the Atom ? electrons ?found in electron cloud ?relative charge of -1 ?relative mass of 0.00055 amu ? protons ?found in nucleus ?relative charge of +1 ?relative mass of 1.0073 amu ? neutrons ?found in nucleus ?neutral charge ?relative mass of 1.0087 amu 12 J.J.Thomson and Cathode Rays ? Cathode rays originated in the negative electrode (cathode) and move to the positive electrode (anode) ? Cathode rays consists of negative charged particles ? electrons ? Thompson determined the charge ?to-mass ratio of an electron ( The Noble Prize in physics in 1906) 13 Plum Pudding Atom ? Work done by J. J. Thomson proved that the atom had pieces called electrons. ? Thomson found that electrons are much smaller than atoms and carry a negative charge. ? Thomson postulated that an atom consisted of a diffuse cloud of positive charge with the negative electrons embedded randomly in it 14 Thompson?s Plum Pudding Model ? ? ?? ? ?? ? ?? ? ? ? ? ? ? ? ? ? ? ? ? 15 Millikan?s Experiment 16 Millikan?s Experiment ? Determine the charge on the electron ? Calculate the mass of the electron as 9.11x10-30 kg 17 Rutherford?s Experiment 18 Thompson?s and Ratherford?s atomic models ? http://www.mhhe.com/physsci/chemistry/es sentialchemistry/flash/ruther14.swf 19 Rutherford?s Model of the Atom ?atom is composed mainly of empty space ?all the positive charge and most of the mass is in a small area called the nucleus ?electrons are in the electron cloud surrounding the nucleus 20 Relative size of atom and atomic nucleus 21 The Nature of Electrical Charge ? Electrical charge is a fundamental property of protons and electrons. ? Positively and negatively charged objects attract each other. ? Like charged objects repel each other. ? + to +, or ? to ?. ? When a proton and electron are paired, the result is a neutral charge. ? Because they have equal amounts of charge. 22 The Nature of Electrical Charge Electrical charges of the same type repel one another, and charges of the opposite type attract one another. 23 Suba tomic parti cl e Mass g Mass amu Location in at om Charge Symbo l Proton 1.672 62 x 10 - 24 1.007 3 nucle us 1+ p, p + , H + Ele ct ron 0.000 91 x 10 - 24 0.000 55 em pty s pace 1 ? e, e - Neutr on 1.674 93 x 10 - 24 1.008 7 nucle us 0 n, n 0 24 Practice?An Atom Has 20 Protons. Determine if Each of the Following Statements Is True or False? ? If it is a neutral atom, it will have 20 electrons. ? If it also has 20 neutrons, its mass will be approximately 40 amu. ? If it has 18 electrons, it will have a net 2? charge. True True False 25 The Periodic Table of Elements Atomic number Atomic mass Element symbol 26 Atomic number, Z ? the number of protons in the nucleus ? the number of electrons in a neutral atom ? the integer on the periodic table for each element 27 Mass Number, A ? integer representing the approximate mass of an atom ? equal to the sum of the number of protons and neutrons in the nucleus 28 Isotopes ? atoms of the same element which differ in the number of neutrons in the nucleus ? designated by mass number 29 SYMBOLS FOR ISOTOPES ? Isotopes are represented by the symbol where Z is the atomic number, A is the mass number and E is the elemental symbol. ? An example of an isotope symbol is . This symbol represents an isotope of nickel that contains 28 protons and 32 neutrons in the nucleus. ? Isotopes are also represented by the notation: Name-A, where Name is the name of the element and A is the mass number of the isotope. ? An example of this isotope notation is magnesium-26. Mg-26 This represents an isotope of magnesium that has a mass number of 26. Ni 6028 30 SYMBOLS FOR ISOTOPES XAZ X - Atomic symbol of the element Z = The Atomic Number, the Number of Protons in the Nucleus (All atoms of the same element have the same no. of protons.) A = The Mass number; A = Z + N N = The Number of Neutrons in the Nucleus N = A - Z Isotopes = atoms of an element with the same number of protons, but different numbers of neutrons in the Nucleus Symbol of the Isotope 31 Isotopes of Hydrogen H-1, 1H, protium ? 1 proton and no neutrons in nucleus ? only isotope of any element containing no neutrons in the nucleus ? most common isotope of hydrogen 32 Isotopes of Hydrogen H-2 or D, 2H, deuterium ? 1 proton and 1 neutron in nucleus 33 Isotopes of Hydrogen H-3 or T, 3H, tritium ? 1 proton and 2 neutrons in nucleus 34 35 Neon 9.25%221210Ne-22 or 0.27%211110Ne-21 or 90.48%201010Ne-20 or Percent natural abundance A, mass number Number of neutrons Number of protonsSymbol Ne2010 Ne2110 Ne2210 36 Practice?Complete the Following Table. At om ic N u m b e r Mass N u m b e r Nu m b e r of p r oton s Nu m b e r of e lec tr on s Nu m b e r of n e u tr on s Calciu m - 40 Car b on - 13 Alu m in u m - 27 +3 37 Practice?Complete the Following Table, Continued. At om ic N u m b e r M ass N u m b e r Num b e r of p r ot on s Num b e r of e lec t r on s Num b e r of n e u t r on s Calc iu m - 40 20 40 20 20 20 Car b on - 13 6 13 6 6 7 Alu m in u m - 27 +3 13 27 13 10 14 PERIODIC LAW ? This is a statement about the behavior of the elements when they are arranged in a specific order. ? In its present form the statement is: Elements with similar chemical properties occur at regular (periodic) intervals when the elements are arranged in order of increasing atomic numbers. 38 Modern Periodic Table Moseley, Henry Gwyn Jeffreys 1887?1915, English physicist. ? studied the relations among spectra of different elements. ? concluded that the atomic number is equal to the charge on the nucleus based on the x-ray spectra emitted by the element. ? explained discrepancies in Mendeleev?s Periodic Law. In Modern Periodic Table the elements are arranged according to increasing atomic numbers 39 Organization of Periodic Table ? period ? horizontal row ? group ? vertical column 40 PERIODIC TABLE ? A periodic table is a tabular arrangement of the elements based on the periodic law. ? In a modern periodic table, elements with similar chemical properties are found in vertical columns called groups or families. group/family period 41 PERIODIC TABLE GROUP OR FAMILY ? A vertical column of elements that have similar chemical properties. ? Traditionally designated by a Roman numeral and a letter (either A or B) at the top of the column. ? Designated only by a number from 1 to 18 in a modern but as yet not universally-used designation. 42 APPEARANCE OF A MODERN PERIODIC TABLE ? In a modern table, elements 58-71 and 90-103 are not placed in their correct periods, but are located below the main table. 43 I A I I A I I I B I V B V B V I B V I I B V I I I B I B I I B I I I A I V A V A V I A V I I A V I I I A 1 1 2 1 H H He 1 . 0 0 8 1 . 0 0 8 4 . 0 0 2 6 3 4 5 6 7 8 9 10 2 Li Be B C N O F Ne 6 . 9 3 9 9 . 0 1 2 2 1 0 . 8 1 1 1 2 . 0 1 1 1 4 . 0 0 7 1 5 . 9 9 9 1 8 . 9 9 8 2 0 . 1 8 3 11 12 13 14 15 16 17 18 3 Na Mg Al Si P S Cl Ar 2 2 . 9 9 2 4 . 3 1 2 2 6 . 9 8 2 2 8 . 0 8 6 3 0 . 9 7 4 3 2 . 0 6 4 3 5 . 4 5 3 3 9 . 9 4 8 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 3 9 . 1 0 2 4 0 . 0 8 4 4 . 9 5 6 4 7 . 8 9 5 0 . 9 4 2 5 1 . 9 9 6 5 4 . 9 3 8 5 5 . 8 4 7 5 8 . 9 3 2 5 8 . 7 1 6 3 . 5 4 6 5 . 3 7 6 9 . 7 2 7 2 . 5 9 7 4 . 9 2 2 7 8 . 9 6 7 9 . 9 0 9 8 3 . 8 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 8 5 . 4 6 8 8 7 . 6 2 8 8 . 9 0 6 9 1 . 2 2 4 9 2 . 9 0 6 9 5 . 9 4 * 9 8 1 0 1 . 0 7 1 0 2 . 9 1 1 0 6 . 4 2 1 0 7 . 9 1 1 2 . 4 1 1 1 4 . 8 2 1 1 8 . 7 1 1 2 1 . 7 5 1 2 7 . 6 1 1 2 6 . 9 1 3 1 . 2 9 55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 6 Cs Ba * * L a Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 1 3 2 . 9 1 1 3 7 . 3 3 1 3 8 . 9 1 1 7 8 . 4 9 1 8 0 . 9 5 1 8 3 . 8 5 1 8 6 . 2 1 1 9 0 . 2 1 9 2 . 2 2 1 9 5 . 0 8 1 9 6 . 9 7 2 0 0 . 2 9 2 0 4 . 3 8 2 0 7 . 2 2 0 8 . 9 8 * 2 0 9 * 2 1 0 * 2 2 2 87 88 89 104 105 106 107 108 109 110 111 112 114 116 118 7 Fr Ra * * * A c Rf Ha Sg Ns Hs Mt U u n U u u U n b U u q U u h U u o * 2 2 3 2 2 6 . 0 3 2 2 7 . 0 3 * 2 6 1 * 2 6 2 * 2 6 3 * 2 6 2 * 2 6 5 * 2 6 8 * 2 6 9 * 2 7 2 * 2 7 7 * 2 8 5 * 2 8 9 * 2 9 3 B a s e d o n s y m b o l s u s e d b y A C S S . M . C o n d r e n 2 0 0 1 58 59 60 61 62 63 64 65 66 67 68 69 70 71 * D e s i g n a t e s t h a t * * L a n t h a n u m Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu a l l i s o t o p e s a re Se ri e s 1 4 0 . 1 2 1 4 0 . 9 1 1 4 4 . 2 4 * 1 4 5 1 5 0 . 3 6 1 5 1 . 9 6 1 5 7 . 2 5 1 5 8 . 9 3 1 6 2 . 5 1 1 6 4 . 9 3 1 6 7 . 2 6 1 6 8 . 9 3 1 7 3 . 0 4 1 7 4 . 9 7 ra d i o a ct i v e 90 91 92 93 94 95 96 97 98 99 100 101 102 103 * * * A ct i n i u m Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Se ri e s 2 3 2 . 0 4 2 3 1 . 0 4 2 3 8 . 0 3 2 3 7 . 0 5 * 2 4 4 * 2 4 3 * 2 4 7 * 2 4 7 * 2 5 1 * 2 5 2 * 2 5 7 * 2 5 8 * 2 5 9 * 2 6 0 Periodic Table of the Elements 44 The Modern Periodic Table ? Main group = representative elements = ?A? groups. ? Transition elements = ?B? groups. ? All metals. ? Bottom rows = inner transition elements = rare earth elements. ? Metals ? Really belong in periods 6 and 7. 45 = Alkali metals = Alkali earth metals = Noble gases = Halogens = Lanthanides = Actinides = Transition metals 46 Family Names Group IA alkali metals Group IIA alkaline earth metals Group VIIA halogens Group VIIIA noble gases transition metals inner transition metals ? lanthanum series rare earths ? actinium series trans-uranium series 47 Alkali Metals ? Group IA = Alkali metals. ? Soft, low melting points, low density. ? Flame tests: Li = red, Na = yellow, and K = violet. ? Very reactive, never found uncombined in nature. ? Tend to form water soluble compounds that are crystallized from seawater then molten salt electrolyzed. ? Colorless solutions. ? React with water to form basic (alkaline) solutions and H2: 2 Na + 2 H2O ? 2 NaOH + H2 ? Releases a lot of heat. lithium sodium potassium rubidium cesium 48 Alkaline Earth Metals ? Group IIA = Alkaline earth metals. ? Harder, higher melting, and denser than alkali metals. ? Mg alloys used as structural materials. ? Flame tests: Ca = red, Sr = red, and Ba = yellow-green. ? Reactive, but less than corresponding alkali metal. ? Form stable, insoluble oxides from which they are normally extracted. ? Oxides are basic = alkaline earth. ? Reactivity with water to form H2: Be = none, Mg = steam, Ca, Sr, Ba = cold water. magnesium calcium beryllium strontium barium 49 Halogens ? Group VIIA = Halogens. ? Nonmetals. ? F2 and Cl2 gases, Br2 liquid, and I2 solid. ? All diatomic. ? Very reactive. ? Cl2, and Br2 react slowly with water: Br2 + H2O ? HBr + HOBr ? React with metals to form ionic compounds. ? hydrogen halides all acids: ? HF weak < HCl < HBr < HI. bromine iodine chlorine fluorine 50 Noble Gases ? Group VIIIA = Noble gases. ? All gases at room temperature. ? Very low melting and boiling points. ? Very unreactive, practically inert. ? Very hard to remove electron from or give an electron to. 51 Types of Elements metals nonmetals metalloids ? semimetals 52 Metals ? Solids at room temperature, except Hg. ? Reflective surface. ? Shiny ? Conduct heat. ? Conduct electricity. ? Malleable. ? Can be shaped. ? Ductile. ? Drawn or pulled into wires. ? Lose electrons and form cations in reactions. ? About 75% of the elements are metals. ? Lower left on the table. 53 Nonmetals ? Found in all 3 states. ? Poor conductors of heat. ? Poor conductors of electricity. ? Solids are brittle. ? Gain electrons in reactions to become anions. ? Upper right on the table. ? Except H. 54 Metalloids ? Show some properties of metals and some of nonmetals. ? Also known as semiconductors. Properties of Silicon: ?Shiny ?Conducts electricity ?Does not conduct heat well ?Brittle 55 Ions ?charged single atom ?charged cluster of atoms ? an atom or a group of atoms that has a net positive or negative charge 56 Ions ? cations ?positive ions ? anions ?negative ions Ionic compounds ?combination of cations and anions ?zero net charge 57 58 Ions ? Atoms acquire a charge by gaining or losing electrons. ? Not protons! ? Ion charge = # protons ? # electrons. ? Ions with a positive charge are called cations. ? More protons than electrons. ? Form by losing electrons. ? Ions with a negative charge are called anions. ? More electrons than protons. ? Form by gaining electrons. ? Chemically, ions are much different than the neutral atoms. ? Because they have a different structure. IONIC BOND FORMATION ? Ions with positive charges are attracted to ions with negative charges. The attractive force between such ions holds them together and is called an ionic bond. ? Ionic bonds form when representative metal atoms lose valence electrons. The electrons are gained by representative nonmetal atoms. Both atoms are changed into ions with noble gas configurations. The resulting ions are then attracted to each other. 59 IONIC COMPOUNDS ? The substances formed when ionic bonds form between positive and negative ions are called ionic compounds. ? When ionic compounds are formed by the reaction of only two elements the resulting ionic compound is called a binary ionic compound. Cu2O CuO 60 FORMULAS FOR BINARY IONIC COMPOUNDS ? Binary ionic compounds typically form when a metal and a nonmetal react. ? The metal tends to lose one or more electrons and forms a positive ion. ? The nonmetal tends to gain one or more electrons and forms a negative ion. ? The symbol for the metal is given first in the formula. ? The formula for a binary ionic compound represents the minimum number of each ion that when combined together will provide equal numbers of positive and negative electrical charges. NaCl 61 EXAMPLES OF FORMULAS FOR BINARY IONIC COMPOUNDS ? Sodium and fluorine: ? Sodium, a group IA metal, will form sodium ions with the symbol Na+. ? Fluorine, a group VIIA nonmetal, will form ions with the symbol F-. ? The minimum number of ions needed to give the same number of positive and negative charges is one of each. ? The one Na+ provides one positive charge and the one F- provides one negative charge. ? The correct formula that results is NaF. 62 ? Sodium and sulfur: ?Sodium is a group IA metal and will form sodium ions with the symbol Na+. ?Sulfur is a group VIA nonmetal and will form ions with the symbol S2-. ?The minimum number of ions required to give the same number of positive and negative charges is two Na+ ions and one S2- ion. ?The two Na+ ions provide two positive charges and the one S2- ion provides two negative charges. ?The resulting formula is Na2S. 63 ? Aluminum and oxygen: ?Aluminum is a group IIIA metal and will form ions with the symbol Al3+ . ?Oxygen is a group VIA nonmetal and will form ions with the symbol O2-. ?The minimum number of ions required to give the same number of positive and negative charges is two Al3+ ions and three O2- ions. ?The resulting formulas is Al2O3. 64 NAMING BINARY IONIC COMPOUNDS (Type I) ? Binary ionic compounds are named using the following pattern: ? name = metal name + stem of nonmetal name + -ide ? The stem names and ionic symbols for some common nonmetals are given in the following table: 65 EXAMPLES OF BINARY IONIC COMPOUND NAMES (Type I) ? Name K2O: name = metal name + nonmetal stem + -ide name = potassium + ox- + -ide = potassium oxide ? Name Mg3N2: name = metal name + nonmetal stem + -ide name = magnesium + nitr- + -ide = magnesium nitride ? Name BeS: name = metal name + nonmetal stem + -ide name = beryllium + sulf- + -ide + beryllium sulfide ? Name AlBr3: name = metal name + nonmetal stem + -ide name = aluminum + brom- + -ide name = aluminum bromide 66 NAMING BINARY IONIC COMPOUNDS (Type II) IN WHICH METALS FORM IONS WITH MORE THAN ONE CHARGE ? Some metal atoms, especially those of transition and inner- transition elements form more than one type of charged ion. (e.g. Cobalt forms both Co2+ and Co3+ ions.) ? The binary compounds containing such ions are named following the pattern given earlier with one addition, the number of positive charges on the metal ion is indicated by a Roman numeral in parentheses following the metal name. ? For example, the compounds CoO and Co2O3 contain cobalt ions with 2+ and 3+ charges respectively. Their names are cobalt (II) oxide and cobalt (III) oxide. 67 (metals that have >1 possible oxidation state) 68 POLYATOMIC IONS ? Polyatomic ions are covalently-bonded groups of atoms that carry a net electrical charge. Most common polyatomic ions are negatively charged. ? Polyatomic ions names must be memorized to name the compounds containing them. 69 EXAMPLES OF FORMULAS FOR IONIC COMPOUNDS CONTAINING POLYATOMIC IONS ? Compound containing K+ and ClO3-: KClO3 ? Compound containing Ca2+ and ClO3-: Ca(ClO3)2 ? Compound containing Ca2+ and PO43-: Ca3(PO4)2 70 71 NAMING IONIC COMPOUNDS THAT CONTAIN A POLYATOMIC ANION ? The names of ionic compounds that contain a polyatomic anion are obtained using the following pattern: name = name of metal + name of polyatomic anion ? Examples: ?KClO3 is named potassium chlorate ?Ca(ClO3)2 is named calcium chlorate ?Ca3(PO4)2 is named calcium phosphate ?CaHPO4 is named calcium hydrogen phosphate 72 WRITING FORMULAS OF IONIC COMPOUNDS THAT CONTAIN POLYATOMIC IONS ? The rules for writing formulas for ionic compounds containing polyatomic ions are essentially the same as those used for writing formulas for binary ionic compounds. ? The symbol for the metal is written first, followed by the formula for the negative polyatomic ion. Equal numbers of positive and negative charges must be represented by the formula. ? When more than one polyatomic ion is required in the formula, parentheses are placed around the polyatomic ion before the subscript is inserted. ? ? ? ?3 4 3 4 4 42 3N a P O M g P O N H P O 73 Naming Compounds containing Polyatomic Ions ? NaNO2 Sodium Nitrite BaSO3 Barium Sulfite ? Ca(NO3)2 Calcium Nitrate Na2SO4 Sodium Sulfate ? LiClO4 Lithium Perchlorate ? KClO3 Potassium Chlorate ? RbClO2 Rubidium Chlorite ? CsClO Cesium Hypochlorite 74 COVALENT BONDING ? Covalent bonding is a type of bonding in which the octet rule is satisfied when atoms share valence electrons. The shared electrons are counted in the octet of each atom that shares them as illustrated below for fluorine gas, F2. ? The atoms sharing one or more pairs of electrons are each attracted to the shared electrons, and thus, are attracted to each other. The attraction to each other is called a covalent bond. The covalent bond may be represented by the shared pair or by a single line between the bonded atoms. 75 ? Electron sharing resulting in covalent bonding can occur between identical atoms or between different atoms. ? Molecules such as Cl2, O2 and N2 are formed when electron sharing occurs between identical atoms. ? Molecules such as H2O, and CH4 are formed when electron sharing occurs between different atoms. H2O CH4 76 NAMING BINARY COVALENT COMPOUNDS ? The pattern used to name binary covalent compounds is similar to that used to name binary ionic compounds: name = name of least electronegative element + stem of more electronegative element + -ide ? In addition to the pattern, the number of each type of atom in the molecule is indicated by means of the following Greek prefixes: ? Note: The prefix mono is not used when it appears at the beginning of the name. 77 EXAMPLES OF NAMING BINARY COVALENT COMPOUNDS SO2: name = sulfur + di- + ox + -ide = sulfur dioxide XeF6: name = xenon + hexa- + fluor + -ide = xenon hexafluoride H2O: name = di- + hydrogen + mono- + ox + -ide = dihydrogen monoxide (also known as water) (Note, the final o of mono- was dropped for ease of pronunciation.) 78 Acids 79 80 Acids ? Acids are molecular compounds that form H+ when dissolved in water. ? Sour taste. ? Dissolve many metals. ? Like Zn, Fe, Mg, but not Au, Ag, Pt. ? Formula generally starts with H. ? E.g., HCl, H2SO4. 81 Acids ? Contain H+ cation and anion. ? In aqueous solution. ? Binary acids have H+ cation and nonmetal anion. ? Oxyacids have H+ cation and polyatomic anion. 82 Naming Binary Acids ? Write a hydro- prefix. ? Follow with the nonmetal name. ? Change ending on nonmetal name to ?ic. ? Write the word acid at the end of the name. HCl ? hydrocloric acid HBr ? hydrobromic acid HI ? hydroiodic acid 83 Naming Oxyacids ? If polyatomic ion name ends in ?ate, then change ending to ?ic suffix. ? If polyatomic ion name ends in ?ite, then change ending to ?ous suffix. ? Write word acid at end of all names. 84 Example?Naming Oxyacids, H2SO4 1. Identify the anion. SO4 = SO42- = sulfate. 2. If the anion has ?ate suffix, change it to ?ic. If the anion has ?ite suffix, change it to ?ous. SO42- = sulfate ? sulfuric. 3. Write the name of the anion followed by the word acid. sulfuric acid 85 Formula-to-Name Flowchart Guzel Lecture 2 Chapter 2