Chapter 7: Chemical Formulas and Chemical Compounds
I. Chemical Names and Formulas
A. There are millions of natural and synthetic compounds. Many times, the common names for them do not give any information on the chemical composition. To solve this problem, we give all compounds formulas that give this information.
B. Significance of a Chemical Formula
1. For a molecular compound, the chemical formula gives the number of atoms of each element that is present in each molecule of that compound. For example, Octane, C8H18.
2. In ionic compound, the formula gives the lowest ions of atoms in the compound. For example, Aluminum Sulfate, Al2(SO4)3.
C. Monatomic Ions
1. Monatomic Ions are ions that are formed from a single atom.
2. For all S and P block elements, charge can easily be predicted using group number and the Octet Rule. For example, Na+.
3. For many D block elements, several charges can be formed. Know charges for Cu, Ag, Cr, Fe, Pb, Sn, Zn, and Ni.
4. Naming Monatomic Cations
a. Simply give name of element. For example, K+ is potassium ion.
b. Transition metals with multiple charges are named by giving the name with the charge in Roman numerals in parentheses. For example, Cr2+ would be chromium (II).
5. Naming Monatomic Anions
- Drop the ending of the name and add -ide. For example, F- is the fluoride ion.
6. See Table 1 on page 221.
D. Binary Ionic Compounds
1. Binary Compounds are compounds that are composed of 2 different elements.
2. In a binary ionic compound the total number of positive charges and negative charges must be equal. The formula can be written when given the identities of the compound’s ions.
3. Steps to writing the formulas for binary ionic compounds.
a. Write the symbols and charges for the ions side by side. Write the cation first.
b. Cross over the charges by using the absolute value of each ion’s charge as the subscript for the other ion.
c. Check the subscripts and divide them by their largest common factor to give the smallest possible whole number ratio of ions. d. Write the formula, without the charges.
E. Naming Binary Ionic Compounds
1. The nomenclature of binary ionic compounds is simple. Write the name of the cation (usually a metal), followed by the name of the anion.
2. For most simple ionic compounds, the ratio of the ions is not indicated in the name. The ratio is understood, based on the relative charges of the ions.
3. When dealing with ions with multiple charges (d-block), the name of the cation is written, followed immediately by the Roman numerals for the charge. Then write the name of the anion. This is called the Stock System.
4. Formula Writing and Nomenclature Involving Polyatomic Ions
a. Polyatomic Ions are ions that contain more than 1 atom.
b. Formula writing is no different, just remember that the polyatomic ion is 1 thing, not more than that.
c. Many polyatomic ions contain oxygen. They are called oxyanions.
d. In many cases, there are more than one oxyanions formed by the same two elements. When two oxyanions occur the most common oxyanion is given the ending -ate. The one with 1 less oxygen is given the -ite ending. For example, nitrate and nitrite.
e. When more than two oxyanions occur the one with 1 less oxygen than -ite is given the prefix hypo- in addition to the -ite. The one with 1 more oxygen than the -ate is given the prefix per- in addition to the -ate. For example, hypochlorite, chlorite, chlorate, perchlorate.
f. Know ammonium, acetate, chlorate, chlorite, hypochlorite, perchlorate, cyanide, bicarb, bissulfate, hydroxide, nitrate, nitrite, carbonate, chromate, dichromate, sulfate, sulfite, and phosphate.
g. See Table 2 on page 226.
F. Naming Binary Molecular Compounds
1. For binary molecular compounds we use prefixes to give the number of atoms of each element that is in each molecule.
2. Memorize all 10 prefixes, mon, di, tri, tetra, penta, hexa, hepta, octa, nona, and deca.
3. Steps for naming BMCs:
a. Write the name of the less EN element first. It is given a prefix ONLY if there are more than one atom present.
b. The second element is named by combining (a) a prefix indicating the number of atoms of that element present, (b) the root of the second element, and (c) the ending -ide.
c. The “o” or “a” at the end of a prefix is usually dropped when the word following the prefix begins with a vowel, for example: monoxide, pentoxide.
4. Practice with N2O, NO, NO2, N2O3, N2O4, and N2O5.
G. Covalent-Network Compounds
1. Some covalent molecules do not exist of individual molecules. They instead are part of a 3 dimensional network.
2. When this occurs the lowest ratio is given and then named just as binary covalent compounds are.
H. Acids and Salts
1. An acid is a distinct type of compound which will be discussed later.
2. There are 2 types of acids.
a. Binary acids: consist of two elements, usually H and a halogen.
b. Oxyacids: contain H, O, and a third element, usually a nonmetal.
3. Most acids are aqueous solutions, they are dissolved in water.
4. For binary acids with hydrogen, write hydro- for hydrogen, and then the root for the other element followed by -ic, then the word acid. For example, HCl is called Hydrochloric acid.
5. For oxyacids, we must look at the ions produced. Most will produce H+ ions and 1 other ion.
a. If the other ion ends in -ate, the acid ends in -ic. (ex. HNO3)
b. If the other ion ends in -ite, the acid ends in -ous. (ex. HNO2)
c. If the other ion is hypo-X-ite, the acid is hypo-X-ous. (ex HClO)
d. If the other ion is per-X-ate, the acid is per-X-ic. (ex. HClO4)
6. A salt is an ionic compound composed of a cation and an anion.
7. Salts are named just as other ionic compounds.
II. Oxidation Numbers
A. In order to indicate the general distribution of electrons among the bonded atoms in a molecular compound or polyatomic ion, oxidation numbers, or oxidation states, are assigned to the atoms composing the compound.
B. Assigning Oxidation Numbers
1. The atoms in a pure element have an oxidation number of O. (Ex: Na, O2, P4, S8 are all 0)
2. The more electronegative element in a binary molecular compound is assigned the number equal to the negative charge it would have as an anion. The less electronegative atom is assigned the number equal to the positive charge that it would have as a cation.
3. F has an oxidation number of -1 in all compounds because it is the most electronegative element.
4. O has an oxidation number of -2 in almost all compounds. Exceptions include when it is in peroxides (O22- , Ox # = -1) and when it is in compounds with halogens. (In OF2 Ox # =+2)
5. H has an oxidation number of +1 in all compounds containing elements that are more electronegative than it is. It has an oxidation number of -1 in compounds with metals.
6. The algebraic sum of the oxidation numbers on all atoms in a neutral compound is 0.
7. The algebraic sum of the oxidation numbers on all atoms in a polyatomic ion is equal to the charge of the ion.
8. Although rules 1-7 apply to covalently bonded atoms, oxidation states can be assigned to atoms in ionic compounds.
9. Practice with: UF6, H2SO4, ClO3-, KClO4, SO32-, NF3, CO2, NO3-, NH4Cl. (Write Ox # of each atom in compounds)
III. Using Chemical Formulas
A. Formula Masses
1. A formula mass is the sum of the average atomic masses of all the atoms represented in its formula.
2. The formula mass of a formula tells the mass (in amu) of each molecule of the formula.
3. Practice with H2, H2SO4, and CO2.
B. Molar Masses
1. A molar mass is the mass in g of 1 mole (6.022 * 1023 particles) of the substance.
2. It is calculated just like a formula mass, but the units are g/mol.
3. Formula mass and molar mass will be numerically identical, but the units are different.
4. Practice with O2, H3PO4, and Ba(NO3)2.
C. Molar Mass as a Conversion Factor
1. We can use molar mass as a conversion factor between moles of a substance and grams of a substance.
2. Practice with 2.50 moles of O2, 33 g of ibuprofen (C13H18O2).
D. Percent Composition
1. Percent Composition is the percentage by mass of each element in the compound.
2. To get the % comp, divide the mass of element in 1 mole of compound by the molar mass of the compound, then multiply by 100.
3. Practice with All elements in CaCO3.
IV. Determining Chemical Formulas
A. An empirical formula is the lowest ratio of atoms in a compound. A molecular formula is the exact ratio of atoms in a compound.
B. To calculate the empirical formula from the percent composition, follow the following steps:
1. Assume 100g sample. This allows us to replace all % signs with g.
2. Convert all g to moles. Use the molar mass of the elements.
3. Divide all moles by the smallest number of moles. This will give the smallest whole number ratio of elements.
4. Write the symbols for each element and the subscripts that are from the previous step. Make sure that all subscripts are whole numbers.
5. Practice, practice, practice.
C. Calculating the Molecular Formula
1. A molecular formula is simply a multiple of the empirical formula. The molar mass of the molecular formula will usually be given.
2. Once the empirical formula is determined, calculate the molar mass of it.
3. X(EF)=Molecular Formula. We know the MM of the molecular formula, and calculated the MM of the empirical formula. Divide MF/EF to get X.
4. Multiply all subscripts in the EF by the value of X. This will give the molecular formula.
5. Practice, Practice, Practice.