Chemical Formulas & Naming (Nomenclature) Chapter 7 Chemical Formulas & Naming (Nomenclature)

Chemical Formulas Chemical formulas indicate the relative # of each type of atom in a compound Ex: C8H18 (octane) Al2(SO4)3 3x everything in ( ) 2 Aluminum 1 Sulfur & 4 Oxygen 8 Carbons 18 Hydrogens

Monatomic Ions Monatomic Ions are formed from a single atom (+) positive ions = cations = lose e- (-) negative ions = anions = gain e- Atoms gain or lose electrons in an attempt to achieve a full valence shell (octet)

Naming Monatomic Ions Cations(+) K+ : potassium Mg2+ : magnesium D-Block Cations Cu+ : copper(I) Fe3+ : iron(III) * Use roman numerals to indicate the positive charge Anions(-) F- : fluoride N3- : nitride O2- : oxide * Element name ending in ide

Binary Ionic Compounds Binary ionic compounds are made from the combination of cations and anions The charges of the positive and negative ions must balance out to zero (neutral). Ex: Mg2+ + Br -  MgBr2

Crossing Over Method Write the cation followed by the anion 2) Cross over charges to form subscripts 3) Check to make sure charges are balanced

Naming Binary Ionic Compounds Combine the names of both ions Ex: Al2O3 = aluminum oxide AgCl = silver chloride CaBr2 = calcium bromide NaCl = sodium chloride

Stock System of Nomenclature Use Roman #’s to indicate the charge of the cation(+) Ex: CuCl2 = copper (II) chloride FeO = iron (II) oxide Fe2O3 = iron (III) oxide SnF4 = tin (IV) fluoride ZnBr2 = zinc (II) bromide

Polyatomic Ions Polyatomic ions are molecules with an overall positive or negative charge Oxyanions- polyatomic ions containing oxygen Ex: NO2- = nitrite NO3- = nitrate ClO2- = chlorite ClO- = hypochlorite ClO3- = chlorate ClO4- = hyperchlorate

Naming Binary Molecular Compounds (covalent bonds) Use prefixes to indicate the # of each type of element in the molecule Write the element with the smaller group number 1st  If both elements have the same group number, write the element with the greater period number 1st Use a prefix is there is more than 1 of the 1st element

Naming Binary Molecular Compounds (covalent bonds) continued… 3) Write the second element always using a prefix to indicate the # of atoms present Note: When an element begins with a vowel, the “o” or the “a” at the end of a prefix is dropped. Example: monoxide, not monooxide

Prefixes 1 = mono 2 = di 3 = tri 4 = tetra 5 = penta 6 = hexa 7 = hepta 8 = octa 9 = nona 10 = deca

Molecular Compounds: Ex: P4O10 = tetraphosphorus decaoxide N2O = dinitrogen monoxide SO3 = sulfur trioxide

Acids & Salts Binary Acids: contain H & a halogen(17) Ex: HCl = hydrochloric acid Oxyacids: contain H, O, & a 3rd element Ex: H2SO4 = sulfuric acid HNO3 = nitric acid Salts: formed from a cation and an anion of an acid Ex: HCl + NaOH  NaCl + H2O (Acid) (Base) (Salt)

Oxidation #’s (states) Indicate the general distribution of electrons among the bonded atoms in a molecular compound or polyatomic ion Rules: The oxidation number of a free (single) element is always zero (Ex. O2, Na) The more EN element has a – charge and the less EN element has a + charge Fluorine (F) always has an oxidation # of -1 (most EN) Hydrogen (H) has an oxidation # of +1 when combined with nonmetals, -1 when combined with metals The oxidation # of oxygen (O) is usually -2 The oxidation # of other elements is based on the elements group # The sum of the oxidation numbers of all of the atoms in a neutral compound is 0 The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion

Examples UF6 H2SO4 ClO3-

Chemical Formulas Review Formula Mass- sum of the average atomic masses of all elements in a compound, units (amu) Ex: H2O Molar Mass- mass of one mole of a substance, numerically equivalent to formula mass, different units (g/mol) Ex: H2SO4

***The Mole Road*** To convert from Moles  Mass multiply (x) multiply (x) # of Atoms Molar Mass(g/mol) Mass(g) Moles 6.02 x 1023 divide (/) divide (/) To convert from Moles  Mass Multiply the moles by the molar mass of the compound To convert from Moles  # of Atoms Multiply the moles by 6.02 x 1023

% Composition The percentage by mass of each element in a compound Mass of Element -------------------------- x 100 = % Composition Mass of Compound Ex: H2SO4

Deriving Chemical Formulas Empirical Formula- the symbols of the elements in a compound w/ subscripts showing the smallest whole # ratio between the atoms Ex: C6H12O6  CH2O Chemical Formula Empirical Formula

Determining Empirical Formulas Assume 100.0 g sample mass (if no sample mass is given) 2) Convert % Composition to Mass for each element 71 % A = 71 g of “A” 29 % B = 29 g of “B”

Determining Empirical Formulas 3) Determine the # of Moles of each element using the Molar Mass 71.0 g A x 1 mol A/10.0 g A = 7.10 mol A 29.0 g of B x 1 mol B/2.01 g B = 14.5 mol B 4) Determine Mole Ratio by dividing both values from step 3 by the lower # A: 7.10 / 7.10 = 1.0 B: 14.5 / 7.10 = 2.0 = AB2

*** If the mole ratio is NOT a whole # ratio after dividing then… Multiply both #’s by the same factor to get a whole # ratio Ex: A = 1.5 x 2 = 3.0 B = 1.0 x 2 = 2.0 A = 2.25 x 4 = 9.0 B = 1.0 x 4 = 4.0

Determining Empirical Formulas Convert % of each Element to Grams(g) Convert Grams(g) to Moles(mol) Divide Moles by Lowest # Determine Whole # Ratio

Practice Empirical Formulas 32.38% Sodium(Na) 22.65% Sulfur(S) 44.97% Oxygen(O) 2) 47.3% Carbon 10.6% Hydrogen 42.1% Oxygen

Calculating Molecular Formulas Divide: Actual Mass / Empirical Formula Mass 2) Multiply: Multiply the # from step 1 by the subscript of each element in the empirical formula Ex: BH3 Actual Mass = 27.67 g/mol

END OF CHAPTER 7 NOTES!