SAT Chemistry Chemical Formulas - Names and Formulas Of Common Acids and Bases

SAT Chemistry Chemical Formulas - Names and Formulas Of Common Acids and Bases

The definition of an acid and a base is expanded later in a first-year chemistry course. For now, common acids are aqueous solutions of hydrogen compounds that contain hydrogen ions, H+. Common bases are aqueous solutions containing hydroxide ions, OH-.
A binary acid is named by placing the prefix hydro- in front of the stem or full name of the nonmetallic element, and adding the ending -ic. Examples are hydrochloric acid (HCl) and hydrosulfuric acid (H2S).
A ternary compound consists of three elements, usually an element and a polyatomic ion. To name the compound, you merely name each component in the order of positive first and negative second.
Ternary acids usually contain hydrogen, a nonmetal, and oxygen. Because the amount of oxygen often varies, the name of the most common form of the acid in the series consists of merely the stem of the nonmetal with the ending -ic. The acid containing one less atom of oxygen than the most common acid is designated by the ending -ous. The name of the acid containing one more atom of oxygen than the most common acid has the prefix per- and the ending -ic, that of the acid containing one less atom of oxygen than the -ous acid has the prefix hypo- and the ending -ous. This is evident in Table 7 with the acids containing H, Cl, and O.
You can remember the names of the common acids and their salts by learning the follow-ing simple rules:


When the name of the ternary acid has the prefix hypo- or per-, that prefix is retained in the name of the salt (hypochlorous acid = sodium hypochlorite).
The names and formulas of some comon acids and bases are listed in Table 7.


As you have seen, a chemical formula is an indication of the makeup of a compound in terms of the kinds of atoms and their relative numbers. It also has some quantitative applications. By using the atomic masses assigned to the elements, we can find the formula mass of a compound. If we are sure that the formula represents the actual makeup of one molecule of the substance, the term molecular mass may be used as well. In some cases the formula represents an ionic lattice and no discrete molecule exists, as in the case of table salt, NaCl, or the formula merely represents the simplest ratio of the combined substances and not a molecule of the substance. For example, CH2 is the simplest ratio of carbon and hydrogen united to form the actual compound ethylene, C2H4. This simplest ratio formula is called the empirical formula, and the actual formula is the true formula. The formula mass is deter¬mined by multiplying the atomic mass units (as a whole number) by the subscript for that element and then adding these values for all the elements in the formula. For example: Ca(OH)2 (one calcium amu + two hydrogen and two oxygen amu = formula mass).
In Chapter 6, the concept of a mole is introduced. If you have 6.02 x 1023 atoms of an element, then the atomic mass units can be expressed in grams, and then the formula mass can be called the molar mass. This is further explained on page 163. Another example is Fe2O3.
It is sometimes useful to know what percent of the total weight of a compound is made up of a particular element. This is called finding the percentage composition. The simple formula for this is:
Another Example
Find the percent compositions of Cu and H2O in the compound CuSO4 • 5H2O (the dot is read “with”).
When you are given the percentage of each element in a compound, you can find the empirical formula as shown with the following example:
Given that a compound is composed of 60.0% Mg and 40.0% O, find the empirical formula of the compound.

1. It is easiest to think of 100 mass units of this compound. In this case, the 100 mass units are composed of 60. amu of Mg and 40. amu of O. Because you know that 1 unit of Mg is 24 amu (from its atomic mass) and, likewise, 1 unit of O is 16, you can divide 60 by 24 to find the number of units of Mg in the compound and divide 40. by 16 to find the number of units of O in the compound.
2. Now, because we know formulas are made up of whole-number units of the elements, which are expressed as subscripts, we must manipulate these numbers to get whole numbers. This is usually accomplished by dividing these numbers by the smallest quo­tient. In this case they are equal, so we divide by 2.5.
Another Example ___________________________________
Given: Ba = 58.81%, S = 13.73%, and O = 27.46%. Find the empirical formula.
In some cases you may be given the true formula mass of the compound. To check if your empirical formula is correct, add up the formula mass of the empirical formula and compare it with the given formula mass. If it is not the same, multiply the empirical formula by the small whole number that gives you the correct formula mass. For example, if your empirical formula is CH2 (which has a formula mass of 14) and the true formula mass is given as 28, you can see that you must double the empirical formula by doubling all the subscripts. The true formula is C2H4.


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