7-1 Chemical Names and Formulas I. Significance of a Chemical Formula: A. Molecular Formulas: number of atoms of each element in one molecule of a compound C2H6 = ethane (2 carbon atoms, 6 hydrogen atoms) B. Ionic Compounds: represents the simplest whole number ratio of compounds’ cations and anions Al2(SO4)3 = aluminum sulfate (2 aluminum ions, 3 sulfate ions) II. Monatomic Ions: ions formed from a single atom A. Naming Monatomic Ions: 1. Monatomic cations are identified by the element’s name 2. Monatomic anions drop the ending of the element name and add an “–ide” ending III. Binary Ionic Compounds: compounds composed of two different elemental ions A. Writing Formulas for Binary Ionic Compounds: 1. Write the symbols for the ions side by side. ALWAYS write the cation first! 2. Cross over the charges by using the absolute value of each ion’s charge as the subscript for the other ion 3. Check that the subscripts are in smallest whole number ratio B. Naming Binary Ionic Compounds from Their Formulas 1. Name the cation then the anion C. The Stock System of Nomenclature: 1. Roman numerals are used to denote the charge of metals that can form two or more cations 2. The numeral is enclosed in parentheses and placed immediately after the metal name a. Iron(II) and Iron(III), pronounced “iron two” and “iron three” 3. Roman numerals are never used for anions and for metals that form only one ion D. Compounds Containing Polyatomic Ions: 1. Oxyanions: polyatomic anions that contain oxygen 2. Naming a series of similar polyatomic ions ClO ClO2 ClO3 ClO4 Hypochlorite Chlorite Chlorate Perchlorate 3. Naming compounds containing polyatomic ions is the same as for monatomic ions 4. Writing formulas including polyatomic ions: a. Use parentheses when you need MORE THAN one of a polyatomic ion b. Parentheses are NEVER used for monatomic ions, regardless of how many are needed IV. Naming Binary Molecular Compounds: A. Binary Molecular Compounds: covalently bonded molecules with two elements, both nonmetals B. Naming: 1. Least electronegative element is named first 2. First element gets a prefix if there is more than 1 atom of that element 3. Second element ALWAYS gets a prefix, and an “-ide” ending Examples: N2O3 = dinitrogen trioxide CO = carbon monoxide, not monocarbon monoxide V. Covalent Network Compounds: like ionic compounds in that there are no distinct units (large 3-D framework) but are named using the same system as binary molecular compounds VI. Acids and Salts: A. Binary Acids: acids that consist of two elements, usually hydrogen and one of the halogens B. Oxyacids: acids that contain hydrogen, oxygen, and a third element (usually a nonmetal) C. Salts: an ionic compound composed of a cation and the anion from an acid D. Naming Binary Acids: the name of the binary acid consists of two words: 1) The first word has three parts: - “hydro” prefix - root of the nonmetal element - “ic” ending 2) The second word is always “acid” Examples: HCl = hydro chlor ic acid = hydrochloric acid HBr = hydro brom ic acid = hydrobromic acid HF = hydro fluor ic acid = hydrofluoric acid E. Naming Oxyacids: these are more difficult to name because these acids have hydrogen, a nonmetal, and may have varying numbers of oxygen atoms. For example, H2SO5, H2SO4, H2SO3, and H2SO2 are all acids. The “ate” ions (sulfate, nitrate, etc) make the “ic” acids (sulfuric acid, nitric acid) Examples: SO42- = sulfate ion H2SO4 = sulfuric acid NO3- = nitrate ion HNO3 = nitric acid The acid with one less oxygen then the -ic acid is called the ___________-ous acid. Examples: H2SO4 = sulfuric acid HNO3 = nitric acid H2SO3 = sulfurous acid HNO2 = nitrous acid 7-2 Oxidation Numbers I. Oxidation Numbers (oxidation states): numbers assigned to atoms composing a compound or ion that indicate the general distribution of electrons among bonded atoms *oxidation numbers give atoms in a covalent compound imaginary charges as if they were ionic *this is useful in naming compounds, writing formulas, and balancing chemical equations II. Using Oxidation Numbers for Formulas and Names: -Be aware that the stock number (Roman numerals) system may be used for molecular compounds as well as ionic compounds. However, we will use the traditional (prefix) method of naming binary molecular compounds. -example: it is known that sulfur has two common oxidation states: +4 & +6. What would be the formulas and names that it would have with oxygen? 7-3 Using Chemical Formulas I. Formula Mass: sum of the average atomic masses of all the atoms represented in the formula of a molecule, formula unit, or ion Formula Mass of glucose, C6H12O6 : C = 12.01amu x 6 = 72.06 amu H = 1.01amu x 12 = 12.12 amu O = 16.00amu x 6 = 96.00 amu Formula Mass =180.18 amu II. Molar Mass: mass in grams of one mole (6.02x1023 particles) of a substance -compound's molar mass is numerically equal to it formula mass, but expressed in units of grams/mole (g/mol) Molar Mass of glucose, C6H12O6 = 180.18 g/mol -molar mass can be used as a conversion factor to relate an amount in moles to a mass in grams for a given substance and the reciprocal of molar mass can be used to relate a mass in grams to the number of moles of a substance amount in moles x molar mass(g/mol) = mass in grams mass in grams x reciprocal of molar mass (mol/g) = amount in moles example: What is the mass in grams of 2.50mol of H2O? 1) must find molar mass of water: H: 1.01amu x 2 = 2.02 g/mol O: 16.00amu x 1 = 16.00g/mol 18.02g/mol 2) use molar mass to find mass of compound: 2.50 mol H2O x 18.02 g = 45.0 g H2O 1 mol III. Percentage Composition: the percentage by mass of each element in a compound mass of element in 1 mol of compound x 100 = % element in compound molar mass of compound example: Find % composition of glucose, C6H12O6. 6 mol C x 12.01g/mol = 72.06 g C 72.06/180.18 x 100 = 39.99% C 12 mol H x 1.01g/mol = 12.12 g H 12.12/180.18 x 100 = 6.73% H 6 mol O x 16.00g/mol = 96.00 g O 96.00/180.18 x 100 = 53.28% O 7.4 Determining Chemical Formulas Empirical formula: consists of the symbols for the elements combined in a compound, with subscripts showing the smallest whole-number mole ratio of the different atoms -for ionic compounds, empirical formula = formula unit -for covalent compounds, the empirical formula does not necessarily = the actual molecular formula -example: diborane gas: empirical formula = BH3 molecular formula = B2H6 Calculation of Empirical Formulas: Convert % composition to a mass composition by assuming you have 100.0g sample of compound Convert mass composition to mole composition by dividing by the appropriate molar mass Obtain the smallest mole ratio possible by dividing each mole composition by the smallest - If mass composition is given instead of % comp, simply start at step #3 Example: Find the empirical formula of a compound that contains 53.70% Fe and 46.30% S. Calculation of Molecular Formulas: -the relationship between a compound’s empirical formula and molecular formula is: x(empirical formula) = molecular formula -the same is true about empirical formula mass and molecular formula mass: x(empirical formula mass) = molecular formula mass * x is the factor by which the subscripts in the empirical formula must be multiplied to obtain the molecular formula * to determine molecular formula, you must know the compound’s formula mass -dividing the experimental formula mass by the empirical formula mass gives the value of x Example: The molar mass of a compound is 92g/mol. Analysis of a sample of the compound indicates that it contains 0.606 g N and 1.390 g O. Find its molecular formula.