Writing Formulas for Binary Ionic Compounds: ... among the bonded atoms in a molecular compound or a ... 1mole C6H12O6 has 6 moles C, 12 moles. H, and...
7.1 Chemical Names and Formulas • Formulas tell the number and kinds of atoms in a compound. • subscripts tell the number of each atom or group of atoms.
7.1 Chemical Names and Formulas Naming Binary Ionic Compounds:
• contain only 2 elements • monatomic ion: one atom with a charge • 1 positive ion (cation) and 1 negative ion (anion) • name the cation first - keeps same name as the element • name the anion last changing the ending to –ide
7.1 Chemical Names and Formulas Naming Binary Ionic Compounds: • some transition metals may have more than one possible charge • Stock naming system: a Roman numeral in parentheses is placed after the name of the element to indicate the numerical value of the charge.
7.1 Chemical Names and Formulas Writing Formulas for Binary Ionic Compounds: • write the symbol for the cation first • write the symbol for the anion last • add subscripts to balance out the charges – all compounds are neutral! • may use the “criss-cross” method • formulas should NOT have any charges
7.1 Chemical Names and Formulas
Polyatomic Ions • ions made of more than one atom • (most are negative)
7.1 Chemical Names and Formulas Writing Formulas for Compounds with Polyatomic Ions • write the symbol for the cation first, which may be polyatomic (ex.: NH4+) • write the symbol for the anion last, which may be polyatomic (ex.: OH¯) • add subscripts to balance out the charges – all compounds are neutral!
7.1 Chemical Names and Formulas Naming Compounds with Polyatomic Ions •name the cation first •name the anion last •do NOT change the ending of polyatomic ions
7.1 Chemical Names and Formulas Naming Binary Molecular Compounds • write less electronegative atom first • C P N H S I Br Cl O F • add a prefix only if more than one atom • second atom always has numerical prefix and -ide suffix • “o” or “a” on a prefix is dropped when element begins with a vowel
7.1 Chemical Names and Formulas
Writing Binary Molecular Compound Formulas • use the prefixes in the name to tell you the subscript of each element in the formula • write the correct symbols for the two elements with the appropriate subscripts
7.1 Chemical Names and Formulas Naming Acids & Bases • acids produce H3O+ (H+) when dissolved in water • bases produce OH‾ when dissolved in water
7.1 Chemical Names and Formulas
Writing Formulas for Acids & Bases • for acids: use the rules for writing the names of acids in reverse • for bases: write the formula as you would any other ionic compound
7.2 Oxidation Numbers • The charges on the ions in an ionic compound reflect the electron distribution of the compound. • In order to indicate the general distribution of electrons among the bonded atoms in a molecular compound or a polyatomic ion, oxidation numbers are assigned to the atoms composing the compound or ion. • Unlike ionic charges, oxidation numbers do not have an exact physical meaning: rather, they serve as useful “bookkeeping” devices to help keep track of electrons. • In general when assigning oxidation numbers, shared electrons are assumed to “belong” to the more electronegative atom in each bond.
7.2 Oxidation Numbers 1.The atoms in a pure element have an oxidation number of zero. 2.The more-electronegative element in a binary compound is assigned a negative number equal to the charge it would have as an anion. Likewise for the less-electronegative element. 3.Fluorine has an oxidation number of – 1 in all of its compounds because it is the most electronegative element. 4.Oxygen usually has an oxidation number of – 2. Exceptions: In peroxides (like H2O2) oxygen’s ox. # is –1 In compounds with fluorine (like OF2) oxygen’s ox. # is +2.
7.2 Oxidation Numbers 5.Hydrogen has an oxidation number of +1 in all compounds containing elements that are more electronegative than it; it has an oxidation number of – 1 with metals. 6.The algebraic sum of the oxidation numbers of all atoms in an neutral compound is equal to zero. 7.The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion.
7.3 Using Chemical Formulas • Formula Mass: the sum of the average atomic mass units of all the atoms/ions in the formula of a compound expressed in units of amu • Molar Mass: the sum of the average atomic masses of all the atoms/ions in the formula of a compound expressed in units of grams • subscripts also give the mole ratio of atoms in a compound • 1mole C6H12O6 has 6 moles C, 12 moles H, and 6 moles O
7.3 Using Chemical Formulas • molar mass is used to convert between the mass of a substance and the moles of a substance • in the conversion factor the unit mole always has a value of one • 1 mol = molar mass (grams) of a substance
7.3 Using Chemical Formulas
Molar Mass of Gases • Avogadro’s hypothesis: equal volumes of gases at the same temperature and pressure contain equal numbers of particles • standard temperature and pressure (STP) = 0 °C and 1 atm • at STP, one mole of ANY gas occupies a volume of 22.4 L (1mol = 22.4 L)
7.3 Using Chemical Formulas
7.3 Using Chemical Formulas Molarity • The concentration of a solution is a measure of the amount of solute that is dissolved in a given quantity of solvent. • A dilute solution is one that contains a small amount of solute. • A concentrated solution contains a large amount of solute. • Molarity (M) is the number of moles of solute dissolved in one liter of solution. • To calculate the molarity of a solution, divide the moles of solute by the total volume of the solution.
7.3 Using Chemical Formulas Molarity A. To make a 0.5-molar (0.5M) solution, first add 0.5 mol of solute to a 1-L volumetric flask half filled with distilled water. B. Swirl the flask carefully to dissolve the solute. C. Fill the flask with water exactly to the 1-L mark.
7.3 Using Chemical Formulas Making Dilutions • Diluting a solution reduces the number of moles of solute per unit volume, but the total number of moles of solute in solution does not change. • Since the total number of moles of solute remains unchanged upon dilution, you can write this equation:
7.3 Using Chemical Formulas Percent Composition • relative amounts of elements in a compound •
% element =
mass of element x 100% mass of compound
7.3 Using Chemical Formulas % Composition from a Formula • remember subscripts also give the mole ratio of atoms in 1 mole of a compound • % by mass =
mass of element in one mole of compound x 100% molar mass of compound
• percent composition can be used as a conversion factor
7.3 Using Chemical Formulas Hydrates ‣ crystals (ionic compounds) with water molecules attached - contain a specific ratio of water to compound ‣ NiSO3 • 6H2O • = water is attached not multiplied ‣ named: nickel(II) sulfite hexahydrate ‣ NiSO3 is the anhydrous portion
7.4 Determining Chemical Formulas Empirical Formulas • the simplest ratio of elements in a compound • fractions of atoms do not exist therefore elements combine in whole number ratios • the empirical formula is not necessarily the true formula for a molecular compound (it will be for ionic compounds)
7.4 Determining Chemical Formulas Empirical Formulas • to find the empirical formula you must know the moles of each element • if you know the masses you can find the moles • once you know the moles you can find the ratio • divide moles of all elements by the smallest number of moles to get the whole number ratio • must be either .9 or .1 to round to the whole number • hydrate formulas: found the same way as empirical formulas but treat compound and water as two parts of whole
7.4 Determining Chemical Formulas Empirical Formulas if the number ends with close to .5 multiply by 2 if the number ends with close to .3 or .6 multiply by 3 if the number ends with close to .25 or .75 multiply by 4
7.4 Determining Chemical Formulas Molecular Formulas • indicate the actual (true) ratio of atoms in a compound • are either the same as its experimentally determined empirical formula, or are a simple whole-number multiple of its empirical formula • once you know the empirical formula you only need one more piece of information to find the molecular formula and that is the mass of the compound
7.4 Determining Chemical Formulas Molecular Formulas • to calculate the molecular formula you must find the ratio of the mass of the empirical formula to that of the mass of the molecular formula this should be a whole # ratio • (whole # ratio) x (empirical formula subscripts) = molecular formula