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I. Molecules and Ions Chemical Bonds Forces that hold atoms together

I. Molecules and Ions Chemical Bonds Forces that hold atoms together Molecules = collections of atoms that form a new substance (compound) Covalent Bond = sharing of electron pairs between two atoms Chemical Formula = representation of a molecule showing the number and type of atoms

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I. Molecules and Ions Chemical Bonds Forces that hold atoms together

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  1. I. Molecules and Ions • Chemical Bonds • Forces that hold atoms together • Molecules = collections of atoms that form a new substance (compound) • Covalent Bond = sharing of electron pairs between two atoms • Chemical Formula = representation of a molecule showing the number and type of atoms • CO2 has one carbon atom and two oxygen atoms • NH3 has one nitrogen atom and three hydrogen atoms • Structural Formula = representation of a molecule showing individual chemical bonds as lines • Normal lines are taken to be in the plane of the page • Dashed lines are taken to go back behind the plane of the page • Wedged line are taken to go forward in front of the plane of the page

  2. Ionic Bond = attraction between oppositely charged ions • Ion = charged atom of group of atoms NH4+ or NO3- (polyatomic ions) • Cation = ion with a positive charge Na - e- -------> Na+ • Anion = ion with a negative charge Cl + e- -------> Cl- • Such a compound (like NaCl) is called a Salt or an Ionic Solid C Cl-

  3. II. Nomenclature • The need for a system • Early chemists named new compounds however they wished • Examples: quicklime, milk of magnesia, laughing gas • A system for naming makes it easier for chemists to communicate about the more than 5 million known chemical compounds • Type I Binary Ionic Compounds • 1. Binary ionic compounds are composed of 2 elements only (cation + anion) • Salts where the cation can have only one possible oxidation state are type I. These are H+ and the Group I-III metals • Rules • Cation comes first and anion second • Cation keeps the name of its element • Anion changes the element name to an –ide ending • Examples • a. NaCl = sodium chloride b. Li2O = lithium oxide • CaS = calcium sulfide d. MgBr2 = magnesium bromide

  4. C. Type II Binary Ionic Compounds • Cations that can have multiple oxidation states require more specificity. These are usually the transition metal cations. • The charge on the metal ion is given as a Roman Numeral in parentheses • Examples • CuCl = copper(I) chloride b. HgO = mercury(II) oxide • Fe2O3 = iron(III) oxide c. MnO2 = manganese(IV) oxide • Determining the oxidation state of the cation • Use the location of the anion to determine its negative charge • Group 7 = -1, Group 6 = -2, Group 5 = -3 • Assign the appropriate cation charge to make a neutral compound • Examples • Ionic Compounds with Polyatomic Ions • Polyatomic ions have special names that must be memorized • Hg22+ (mercury(I)) and NH4+ (ammonium) are the only common cations • Many common anions are polyatomic • There can be several different oxyanions in a series • –ite is the anion with less oxygens; -ate is the one with more • More than two in a series: hypo- means less; per- means more • Hypochlorite (ClO-), Chlorite (ClO2-), Chlorate (ClO3-), Perchlorate (ClO4-)

  5. Common Polyatomic Ions • E. Hydrated Ionic Compounds • CaSO4 • 0.5H2O calcium sulfate hemihydrate • BaCl2 • 6H2O barium chloride hexahydrate • CuSO4 • 5H2O copper(II) sulfate pentahydrate • mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa- • nona-, deca-

  6. F. Binary Covalent Compounds • Formed between two nonmetals • Rules • Name elements in order as they appear in the formula • Name the first element using its full element name (don’t use –ide) • Name the second element using the –ide ending • Modify the element names with a prefix for how many atoms present • Don’t use mono- for the first element • Examples • N2O = dinitrogen monoxide b. NO = nitrogen monoxide • NO2 = nitrogen dioxide d. N2O3 = dinitrogen trioxide • G. Formulas from names are even more important at this stage • vanadium(V) fluoride = ? • dioxygen difluoride = ? • rubidium peroxide = ? • gallium oxide = ?

  7. H. Naming Acids • Acid = compound that produces H+ when dissolved in water • Think of them as H+ ions attached to anions • When the anion doesn’t contain O, name as hydro_____ic acid • HCl = hydrochloric acid • HCN = hydrocyanic acid • H2S = hydrosulfuric acid • When the anion does contain O, this is an oxyacid • If the anion would be named –ate, then name the acid –ic acid • H2SO4 = sulfuric acid • H3PO4 = phosphoric acid • HNO3 = nitric acid • If the anion would be named –ite, then name the acid –ous acid • H2SO3 = sulfurous acid • HNO2 = nitrous acid

  8. III. Molar Mass • A. The molar mass is the mass in grams of one mole of a compound • B. The relative weights of molecules can be calculated from atomic masses • 1. Water = H2O = 2(1.008 g/mol) + 16.00 g/mol = 18.02 g/mol • 2. 1 mole of H2O will weigh 18.02 g, therefore the molar mass of H2O is 18.02 g • 3. 1 mole of H2O will contain 16.00 g of oxygen and 2.02 g of hydrogen • C. Calculations with molar mass • 1. Example: Juglone = C10H6O3 Find Molar mass and # moles in 0.0156 g • 2. Example: How many molecules of isopentyl acetate (C7H14O2) and how many carbon atoms are released in one bee sting (1 mg)?

  9. D. Mass Percent • Composition of a substance can be given multiple ways • Molecular formula gives ratio of the number of each element • The Mass Percent gives the composition by mass of each element • Example: Ethanol, C2H6O has a molar mass of 46.07 g • Example: Determine the Mass Percent of C14H20N2SO4.

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