Chapter 6 Chemical Formulas

OBJECTIVES 1. Distinguish between ionic and molecular compounds. 2. Define cation and anion and relate them to metal and nonmetal.

 Molecules and Molecular Compounds 1. Molecules a. the smallest electrically neutral unit of a substance b. Made up of two or more atoms

c. examples: 1. Oxygen in air: O 2 2. water: H 2 O d. diatomic molecule: two of the same atom

: seven of them hydrogen (H 2 ) oxygen (O 2 ) fluorine (F 2 ) bromine (Br 2 ) iodine (I 2 ) nitrogen (N 2 ) chlorine (Cl 2 )

 Molecular compounds 1. made up of molecules 2. have low melting and boiling points 3. Exist as gases or liquids at room temperature

4. composed of two nonmetals 5. examples: CO (carbon monoxide) CO 2 (carbon dioxide)

 Ions and Ionic Compounds 1. Ions a. Atoms or group of atoms that are either + charged or – charged b. Cation: + ion lose electrons c. Anion : - ion, gain electrons : change ending to -ide

d. Examples: 1. Sodium loses 1 electron Na becomes Na Chlorine gains 1 electron Cl becomes Cl -1 chloride

2. Ionic compounds a. Consist of: cation (+), anion (-) metal, nonmetal b. Represented by a formula unit c. Solids, high melting point

OBJECTIVES 3. Distinguish between chemical formulas, molecular formulas, and formula units. 4. Use experimental data to show that a compound obeys the law of definite proportions.

 Chemical formulas 1. shows symbols, subscripts 2. gives the number of atoms 3. example: Ca 3 N 2 Ca = 3 N = 2

Al 2 (SO 4 ) 3 Al = 2 S = 1*3 = 3 O = 4*3 = 12

Molecular Compounds 1. Shows the type and number of atoms in a molecule 2. Does not show shape or how the atoms are arranged 3. NaCl, PH 3

 Models that show arrangement of atoms 1. Structural formula H – N – H H 2. Space-filling CO 2

3. Ball and stick molecular model 4. Perspective drawing

 Formula Units 1. represents ionic compounds 2. lowest whole-number ratio of ions 3. example: a. MgCl 2 ratio is 1:2 b. AlCl 3 ratio is 1:3

 Law of Definite Proportions 1. Any sample of a chemical compound, the elements are always combined in the same proportion by mass ex. H 2 O consists: 2 hydrogen atoms, mass of 2 g 1 oxygen atom, mass of 16 g ratio 1:8

 Law of multiple proportions 1. When two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers 2. Examples: FeS FeS 2

OBJECTIVES 5 Use the periodic table to determine the charge of an ion. 6. Define a polyatomic ion and give the names and formulas of the most common polyatomic ions.

 Ionic Charges A. Monatomic ions 1. ions of one atom 2. examples: Na + or Cl -

B. Transition metals (Groups 3-12) 1. Have more than 1 positive charge 2. Charge is shown with a roman numeral 3. Examples: iron (II) and iron (III)

4. Old names a. Iron (II): ferrous ion b. Iron (III): ferric ion c. lower roman numeral ends in –ous d. Higher roman numeral ends in -ic

 Polyatomic ions 1. Tightly held group of atoms acting like one ion 2. Examples: a. SO 4 -2 b. NH 4 +1

OBJECTIVES: 7. Apply the rules for naming and writing formulas for binary ionic compounds 8. Apply the rules for naming and writing formulas for ternary ionic compounds.

 Ionic Compounds A. Binary compounds 1. has 2 elements 2. ALWAYS metal followed by a nonmetal

B. Formula writing 1. Write the symbol and charge of the first element 2. Write the symbol and charge of the second element 3. Charges must cancel and = 0

4. Examples: a. sodium fluoride Na +1 F =0 NaF b. Calcium iodide Ca +2 I ≠ 0 CaI 2

c. Aluminum sulfide Al +3 S ≠ 0 Al 2 S 3

d. Copper(II) oxide gives the charge of copper Cu +2 O =0 CuO

C. Using polyatomic ions 1. tertinary – contains three different elements 2. must use ( ) if there is more than one 3. examples: a. Sodium hydroxide

b. Calcium phosphate c. Ammonium sulfide

D. Naming binary 1. write the name of the first element 2. write the name of the second element and change the ending to “ide” 3. Transition metals **Must include roman numerals

4. Examples: a. CaF 2 Calcium fluoride b. KBr potassium bromide c. CuCl 2 copper (II) chloride

Objectives 9. Apply the rules for naming and writing formulas for binary molecular compounds.

 Writing covalent compound formulas A. Rules 1. Write the symbol and subscript of the first element.(look at prefix) 2. Write the symbol and subscript of the second element.(look at prefix)

3. prefixes **1-mono 6-hexa (only for 2 nd element) 2- di7-hepta 3-tri8-octa 4-tetra 9-nona 5-penta 10-deca

4. Examples: a. tricarbon difluoride C3F2C3F2 b. heptabromine tetrachloride Br 7 Cl 4

B. Naming 1. Write the name of the element using the prefixes. *** Do not use mono for the 1 st element 2. Write the name of the second element using prefixes, change ending to –ide.

3. Examples: a. P 3 O 5 b. SO 4 c. N 2 O

OBJECTIVES 10. Name and write formulas for common acids.

 Naming and writing acids A. Naming binary hydrides 1. use the prefix hydro- and the suffix –ic 2. HF hydrofluoric acid 3. HBr hydrobromic acid

B. Writing binary hydrides 1. hydrochloric acid HCl 2. hydrosulfuric acid H2SH2S

Naming acid containing oxygen 1. use the suffixes –ous or -ic with the negative ion 2. -ous (lower oxygen state) 3. -ic (higher oxygen state) 4. examples: a. HNO 3 nitric acid b. HNO 2 nitrous acid

c. H 3 PO 4 phosphoric acid d. H 3 PO 3 phosphorous acid ** If there is only one ion use –ic.

Writing acids with oxygen 1. acetic acid H +1 C 2 H 3 O 2 -1 HC 2 H 3 O 2 2. oxalic acid H +1 C 2 O 4 -2 H 2 C 2 O 4

Extra oxygens 1. ClO 4 -1 perchlorate 2. ClO 3 -1 chlorate 3. ClO 2 -1 chlorite 4. ClO -1 hypochlorite

5. sulfurs a. Persulfate SO 5 -2 b. Hyposulfite SO 2 -2

Homework problems 1. Name the following acids a. HC 2 H 3 O 2 b. H I c. HNO 3 d. HBr 2. Write the formulas a. Hydrophosphoric acid b. Hydrofluoric acid c. pernitric acid d. Carbonic acid