Chemical Formulas And Chemical Compounds Chapter 7.

Chemical FormulasAnd Chemical Compounds

Chapter 7

• Many compounds go by common names such as salt (NaCl) or limestone (calcium carbonate.) These names do not give information about the compounds.

• To describe the atomic makeup of compounds, chemists use systematic methods for naming compounds and for writing chemical formulas

Chemical Formula A chemical formula tells us the relative number of each type

of atom in a chemical compound. Formulas have the element followed by a subscript.

Ex: C8H18 Tells us there are 8 carbons and 18 hydrogens.

Ionic and covalent compounds have separate own rules that must be obeyed in writing the formula and naming the compound

Ionic Compounds• We are familiar with determining the charge of monatomic

ions in the s and p blocks.• Monotomic ions: ions formed from a single atom.• Monatomic cations are identified by the element’s name. • Ex: Li+ is simply called lithium.• Naming monatomic anions is slightly more complicated:• Drop the ending of the element’s name• add “–ide” to the root name.

• Examples:• F- fluoride (fluorine – replace “ine” with “ide”)• N-3 nitride • O-2 oxide

Binary Ionic Compounds• Binary Compounds: Compounds composed of two elements.• In a binary ionic compound, the total number of positive and

negative charges must equal zero. In other words, the net charge must be zero, and the overall compound is neutral. (criss-cross rule)

• Example:• Mg2+ + Br- = MgBr2

Naming Ionic Compounds• Nomenclature: the naming system of binary ionic compounds

involves combining the names of the compound’s positive and negative ions.

• How to name binary ionic compounds:• The name of the cation is first, followed by the name of the

anion.• The ratio of ions is not indicated in the compound’s name

because it is understood that the overall charge of the compound is neutral.

• Examples:• BaF2

• CaO• AlBr3

You Practice:• Write the formulas for the binary ionic compounds formed

between the following elements• Potassium and Iodine

• Magnesium and Chlorine

• Sodium and Sulfur

• Aluminum and Sulfur

• Aluminum and nitrogen

• Additional practice: #2 on page 223.

Transition Metals• Some elements, such as iron, form two or more cations with

different charges.• To distinguish the ions formed by these elements, the Stock

system of nomenclature is used.• Roman numerals are used to indicate the charge of the ion.• The roman numeral is enclosed in parentheses and placed immediately after the metal name.

• Examples: • Silver(I) nitrate• Iron (II) sulfate

Transition Metals with Charges

Examples• Name the following compounds: • CuCl2

• VBr3

• Write the formulas of the following compounds:• Iron(III) sulfide

• Vanadium (II) fluoride

Practice:• Name the compound or write the formula for each compound

(from page 225 practice problems):• CuO• CoF3

• SnI4

• K2O

• Copper(II) bromide• Mercury (II) sulfide• Lead (IV) sulfide

• More problems are in your book.

Polyatomic Ions

• Polyatomic ions are groups of atoms that are covalently bonded. They are usually anions.• They have a charge that is spread over the ion.• Ex: nitrite NO2

-

• Ex: nitrate NO3-

Polyatomic Ions

• Compounds containing polyatomic ions are named the same way as binary ionic compounds.• The name of the cation is named first, followed by the

name of the anion, but the name of the polyatomic ion is used for the anion.• Ex: Silver Nitrate and Silver Nitrite• Some other examples:• NaOH

• Sn4+ and SO42-

Practice (p. 227 practice)• Give the formula or name of the compound, whichever is

appropriate (use your polyatomic ion paper!!):• Copper(II) sulfate• Ca(OH)2

• KClO3

• Lithium nitrate• Ag2O

• Sodium carbonate• Fe2(CrO4)3

• Potassium perchlorate

Naming Covalent Compounds

•Naming covalent compounds made up of two elements are named by similar method as ionic compounds• The first element named is usually the

first one written in the formula.• The second element has the ending

“–ide.”

Naming Covalent Compounds

• Because covalent compounds are often made of the same elements, we must distinguish between two molecules made of the same elements.• For example, NO and NO2 cannot have the same name!• We use a system of prefixes to show the number of

atoms of each element in the molecule (see table 3, page 228).• The o or a at the end of a prefix is usually dropped when

the word following the prefix begins with another vowel• For instance, we would write monoxide and pentoxide

instead of mono-oxide and penta-oxide.

Prefixes for Covalent compoundsNumber Prefix

1 Mono-

2 Di-

3 Tri-

4 Tetra-

5 Penta-

6 Hexa-

7 Hepta-

8 Octa-

9 Nona-

10 Deca-

Examples:

•Name the following or write the formula•PF5

•XeF4

•CCl4

•Carbon dioxide•Dinitrogen pentoxide• Sulfur hexafluoride

Practice

• SO3

• ICl3

• PBr5

• As2O5

• P4O10

• Carbon tetraiodide• Phosphorus trichloride• Dinitrogen trioxide• Silicon dioxide• Tetranitrogen pentoxide

Naming Acids• Acid: a distinct type of molecular compound.• We will focus on acids later in the year.• Binary acids: Acids that consist of two elements, usually

hydrogen and one of the halogen (F, Cl, Br, or I)• Oxyacids: Acids that contain hydrogen, oxygen, and a third

element (usually a nonmetal)

Binary Acids• To name binary acids:• Use the prefix “hydro” followed by the second element with

the ending “ic acid”• HCl Hydrochloric acid• HF Hydrofluoric acid• HBr Hydrobromic acid• HI Hydriodic acid

Oxyacids• Most oxyacids involve a polyatomic ion.• For acids that contain a polyatomic ion that ends in “ate,”

change the “ate” to “-ic acid.”• For acids that contain a polyatomic ion that ends in “ite,”

change the “ite” to “–ous acid.”• H2SO4 SO4 is the sulfate ion, so this is called Sulfuric acid

• H2SO3 SO3 is the sulfite ion, so this is called Sulfurous acid

Practice:• Name the following acids or provide the correct formula• HF• Sulfuric acid• CH3COOH

• H2CO3

• Phophoric acid• HI• HNO2

• Perchloric acid• H2SO3

Oxidation Numbers

• Oxidation numbers: indicate the general distribution of electrons among the bonded atoms in a molecular compound or polyatomic ion.• Unlike ionic charges, oxidation numbers do not

have an exact physical meaning.

Assigning Oxidation Numbers1. Atoms in a pure element have an oxidation number of zero.2. The more-electronegative element in a binary compound is

assigned the number equal to the negative charge it would have as the anion. The less electronegative atom is assigned the number equal to the positive charge it would have as the cation.

3. Fluorine always has oxidation number of -1.4. Oxygen has oxidation number of -2 except when in

peroxides (-1) and in compounds with fluorine (+2)5. Hydrogen has oxidation number of +1 when bonded to

more electronegative elements. -1 when bonded with metals.

6. The sum of the oxidation numbers in a neutral compound is zero, and sum is equal to charge of ion for a polyatomic ion.

Example:

• UF6

• Rule 3 tells us the oxidation number of F is always -1.• -1 x 6 = -6• Since UF6 is a neutral compound, positive charge must be

+6.• +6 / 1 = +6 (1 from one U atom)

• H2SO4

• O and S are more electronegative than H, so H has oxidation number of +1 (rule 5.)

• O is not a peroxide or halogen, so O is -2 (rule 4)• 1(2) + -2(4) = 2 – 8 = -6• So sulfurs oxidation state is -6

• ClO3-

• Charge of anion is -1, so sum of oxidation numbers will be -1.• The oxidation number of Oxygen is -2 (rule 4)• -2(3) + Cl = -1• -6 + Cl = -1• Cl = +5

Practice• See page 234. Work through number 1 and complete for

homework.

Percent Composition• Percentage Composition: the percentage by mass of each

element in a compound.

• Find the % composition of Copper(I)sulfide, Cu2S:• Find the molar mass of the compound.• Determine the mass of each element present and multiply by its

subscript.• Divide mass of element by mass of compound.

Practice• Find the percent compositions of:• PbCl2

• Ba(NO3)2

• Magnesium hydroxide is 54.87% oxygen by mass. How many grams of oxygen are in 175 grams of the compound? How many moles of oxygen is this?

Empirical Formulas• An empirical formula shows the simplest ratio of the atoms in

a compound• A molecular formula is a whole number multiple of the

empirical formula.

Compound Empirical Formula

Molecular Formula Molar Mass (g)

Formaldehyde CH20 CH20-Same as empirical formula

30.03

Acetic Acid CH20 C2H402

-2x empirical formula60.06

Glucose CH20 C6H1206

-6x empirical formula180.18

Determining Empirical Formula• If the identities of the elements in a compound

are known, then the empirical formula can be determined from:• From % Composition or Mass:

1. If given % composition, assume 100g sample and change % to grams.

2. Convert grams to moles.3. Divide each number of moles by the smallest

value.4. Use the numbers from step 2 as your subscripts.

Determining Empirical Formula

Example problem• Analysis of a compound shows that it contains 32.38 % Na,

22.65 % S, and 44.99 % O. Find the empirical formula.

Practice:• Analysis of a 10.15 g sample of a compound known to contain

only phosphorous and oxygen indicates a phosphorous content of 4.433 g. What is the empirical formula of this compound.

• P. 247 practice problems.

To determine molecular formula• X(empirical formula) = molecular formula• To determine the molecular formula of a compound, you must

know the compound’s formula mass.• For example: • The empirical formula of a compound is BH3. The mass of this

molecule is 13.84 g.• The mass of the unknown compound is 27.67 g.• X=27.67/13.84 = 2.000• So the molecular formula is 2(BH3) = B2H6

• Try the practice problems on page 249.