Chemical Compounds Chapter 3

Chemical Compounds Chapter 3 Molecular Compounds Molecular Formulas exact number of atoms of each type in compound example - H2O  inorganic compounds    contain no carbon examples - NH3, H2SO4, P2O5 many are ionic organic compounds   contain carbon examples - C2H6, C2H6O, C3H8 generally molecular Molecular Compounds Molecular Formulas methods for writing formulas  molecular formula C2H6 (ethane)  C2H6O (ethanol) structural formula H H H C C H H H H H H C C H H O H Functional group  (hydroxide) condensed formula CH3CH3 CH3CH2OH Naming Binary Inorganic Compounds molecules containing atoms of only two elements  element farthest left first HCl (hydrogen chloride)  if same group - lowest first SO2 second element ends in -ide  Greek prefixes for numbers  sulfur dioxide exception - binary hydrogen compounds example: H2S - hydrogen sulfide Naming Binary Inorganic Compounds Naming Binary Inorganic Compounds compounds with common, nonsystematic names: H2O  NH3  N2H4  NO  N2O  PH3  water ammonia hydrazine nitric oxide nitrous oxide phosphine (text p. 82) Hydrocarbons contain only C and H  simplest type is alkanes  general formula - CnH2n+2 (n = integer)  first 4 have common names - remaining systematic  Hydrocarbons Ions and Ionic Compounds ionic compound - composed of positive and negative ions  typically metal with non-metal   metal loses electrons to form cation non-metal gains electrons to form anion example - NaCl (formed from Na+ and Cl-) Ionic Charges determine charges of monatomic ions from Periodic Table Ionic Compounds formed from a cation(s) and an anion(s)  overall charge on formula is neutral examples: Na+ & ClCa2+ & ClMg2+ & N3- NaCl CaCl2 Mg3N2  “empirical” formulas only  write formula: cation then anion Naming Ionic Compounds Cations  Cations from metals  same name as metal Na+ sodium ion  Same metal, differing charges   use Roman numeral to distinguish usually transition metals Fe2+ iron(II) ion  Ca2+ calcium ion Fe3+ iron(III) ion Cations from non-metals  names end in -ium NH3 ammonia NH4+ ammonium Naming Ionic Compounds Anions  Monatomic  drop ending, add -ide H- hydride ion  O2- oxide ion Polyatomic  oxoanions (if only 2 possible - e.g. N, S, P)   larger # of O atoms - suffix -ate NO3- nitrate ion SO42- sulfate ion smaller # of O atoms - suffix -ite NO2- nitrite ion SO32- sulfite ion Naming Ionic Compounds Anions (cont’d)  Polyatomic (cont’d)  oxoanions (if more than two possible - e.g. Cl, Br, I)  largest # of O atoms - prefix per- + suffix -ate ClO4- perchlorate ion IO4- periodate ion  next largest # of O atoms - suffix -ate ClO3- chlorate ion IO3- iodate ion  third largest # of O atoms - suffix -ite ClO2- chlorite ion IO2- iodite ion  smallest # of O atoms - prefix hypo- + suffix -ite ClO- hypochlorite ion IO- hypoiodite ion Naming Ionic Compounds Anions (cont’d)  Polyatomic (cont’d)  oxoanions with hydrogen  name oxoanion portion according to rules  place “hydrogen” in front of oxoanion name  notice that the charge on the anion has increased (become more positive) by the same amount as the number of H+ ions added HSO4hydrogen sulfate SO42- HSO4- Naming Ionic Compounds Anions (cont’d)  Polyatomic anions with nonstandard names CH3COOCNC2O42OHO22O2MnO4CrO42Cr2O72- acetate ion cyanide ion oxalate ion hydroxide ion peroxide ion superoxide ion permanganate ion chromate ion dichromate ion Naming Ionic Compounds  Cation then anion examples: CaCl2 calcium chloride MgSO4 magnesum sulfate (NH4)2CO3 ammonium carbonate Properties of Ionic Compounds metal + nonmetal  crystalline  hard, brittle  high melting points  high boiling points  electrolytes  Percent Composition composition of any compound expressed by   # atoms of each type per molecule or formula unit mass of each element in a mole of compound % by mass of each element in compound example: C12H22O11 342.299 g/mol %C = 12(12.011 g/mol) x 100% = 42.107% 342.299 g/mol %H = 22(1.0079 g/mol) x 100% = 6.4779% 342.299 g/mol ! 11(15.9994 g/mol) x 100% = 51.4151% %O = 342.299 g/mol ! ! (part/whole) x 100% Empirical Formulas ratio of mol of each element gives subscripts H2O - empirical formula for water 2 atoms H + 1 atom O OR 2 mol H + 1 mol O Mass % of element assume 100 g sample Grams of each element use atomic mass Empirical Formula calc. mol ratio (divide each by smallest # mol) Moles of each element Molecular Formula need:   empirical formula molar mass of molecular formula Empirical Formula Molecular Formula multiply calculate Molar mass of empirical formula # to multiply empirical formula by molar mass of molecular molar mass of empirical