Naming Compounds

1
Unit 6
Chemical Formulas

• Naming Compounds
• A. Empirical vs. Molecular Formulas
• 1. Empirical Formula is the lowest terms
  formula. It is used for metal-nonmetal compounds
  (ionic or salts). ex. CaO
• 2. Molecular Formula is a formula that
  expresses the exact number of atoms of all
  elements in a nonmetal compound (covalent or
  molecular). ex. H2O2

2
B. Naming Ionic Compounds (Salts)

• Binary Salts contain only two elements, a metal
  and a nonmetal.
• Name from formula
• ex. NaCl Sodium Chloride
• Name metal then 1 or 2 syllables of ide
• 
  nonmetal
• Formula from name
• ex. Calcium Carbide write down symbols
• Ca2 C-4 cross
  charges
• Ca4C2 lowest terms Ca2C
  

3
2. Stock System

• Used when the first element, the metal in a salt,
  can have more than one possible charge. A Roman
  Numeral is used to tell the Charge of the metal
  ion.
• ex. Give the formula for Iron(II) Oxide
• Iron(II) Fe2 so Fe2O-2 ? FeO
• ex. Give the name of Fe2O3 uncross the
  subscripts Fe3 O-2 so Iron(III) Oxide
• Try these CuF2 PbS2 Zn3N2
• Copper(II) Fluoride, Lead(IV) Oxide, Zinc Nitride
  Zinc has only one charge, no roman numeral

4
3. Ternary Salts

• Contain three or more elements.
• An ionic compound, one or both ions containing
  more than one element a Polyatomic ion.
• a. Polyatomic ions (Groups) are covalently
  bonded atoms that, as a group, have gained or
  lost electrons and have a charge.
• ex. CN cyanide ion
• .CN.

Although the C and N are covalently bonded, they
have a charge and are part of a Salt.
x x x x x
ex NaCN Na .CN.
x x x x x
5
b. Naming Ternary Salts

• Look up group names in table E.
• Group names end with ate or ite except cyanide,
  NH4 ammonium, OH hydroxide
• There are 2 groups that contain one element
  Hg22 dimercury(I) O22 peroxide
• ex. Zn(NO3)2

Zinc Nitrate
ex. Manganese(IV) Chromate
? Mn2(CrO4)4
? Mn(CrO4)2
Mn4 (CrO4)2-
Notice when crossing charges with a group, place
the group in parenthesis and place new subscript
outside of parenthesis.
6
C. Covalent Compounds nonmetal molecular
compounds

• Molecular formulas used (not in lowest terms).
• The elements can combine in different amounts to
  make different molecules.
• ex. N2O and NO2
• Each must have a different name, can use 1.
  Stock system
• N2O Nitrogen(I) Oxide
• NO2 Nitrogen(IV) Oxide

7
2. Prefix Method of Naming Molecules

• Prefixes are used to tell the amount of each
  element in the compound
• 2 di 3 tri 4 tetra 5
  penta
• 6 hexa 7-hepta 8 octa 10 deca
  
• ex. CO2 carbon dioxide
• CCl4 carbon tetrachloride
• P4O8 teraphosphorous octoxide
• dinitrogen heptoxide N2O7
• Beware of compounds with traditional names!
• Ex. H2O water CO carbon monoxide
• NH3 ammonia (found in Table L)

8
II. Stoichiometry the study of the mathematical
relationships implied by chemical formulas and
equations

• Amounts of atoms
• Subscripts in a formula tell us the number of
  atoms of each element in the compound.
• ex. How many oxygen atoms are in H2SO4?
• 4 Oxygen atoms
• ex. How many total atoms are in H2SO4?
• 214 7 atoms
• Subscripts outside of parenthesis are multiplied.
  Ex. How many O atoms are in Al2(SO4)3? 3 x 4
  12 Oxygen atoms.

9
B. Amount of atoms in Hydrates.

• Hydrates are salts physically combined with water
  in a definite ratio.
• The positive and negative ends of the water
  molecule attract the negative and positive ions
  in the salt and become part of the crystal.
• Written as MN xH2O, where the dot means with
  not times.
• ex. How many Oxygen atoms are in CuSO4 5H2O
• 5H2O 10 Hs and 5 Os so there are
• 45 9 Os in Copper Sulfate
  pentahydrate.

.
.
10
C. Mole Amounts.

• We can not see or measure individual atoms in
  lab. We need a convenient group amount that we
  can measure in lab, the mole!
• 600,000,000,000,000,000,000,000 or 6x1023 of
  something is a mole of that thing. It is such a
  large number that it can only be used to count
  very small things such as atoms or molecules.
• A Formula can represent one mole of a sub- stance
  as well as just one of the substance.
• ex. How many moles of Hydrogen atoms are in
  (NH4)2SO3 4H2O ?

.
4x2
16 moles of H atoms
2x4
11
D. Atomic Mass

• Atomic mass is the average mass of all naturally
  occurring isotopes of an element.
• The mass of an average atom of an element.
• Atomic masses are found in the periodic table
  the Regents requires them rounded to the nearest
  tenth.
• ex. What is the atomic mass of a Carbon atom?
• Carbon atomic mass 12.0111amu or 12.0u
• O 16.0u H 1.0u Cl 35.5u Cu
  63.5

12
E. Formula Mass (or Molecular Mass for nonmetals)

• Formula mass is the sum of all atomic masses of
  all atoms of all elements in a compound.
• No. of atoms atomic
  mass
• ex H2 2 x 1.0u 2.0u
• ex. H2O List all elements in the compound
• H 2 x 1.0u 2.0u
• O 1 x 16.0u 16.0u
• 18.0u
• ex. Fe2(CO3)3

ex. CuSO4 5H2O
249.6u
291.6u
13
F. Gram Formula Mass a formula mass expressed
in grams instead of amus.

• formula mass mass of 1 molecule in amus
• 1g 6 x 1023 u
• gram formula mass mass of 6x1023 molecules
• Gram formula mass is the mass of one mole of the
  substance!
• If you have 18.0g of H2O, you have 1 mole of H2O
  molecules.
• If you have 1 mole of NaCl, you have 58.5g of
  NaCl.
• Gram formula mass is also called gram
• molecular mass (for nonmetals) or Molar Mass

14
G. Mole Mass Conversions

• How many moles of water molecules are in a 36.0g
  sample of water?
• We know 1 mole of water weighs 18.0g.
• If we divide the given mass by this gram formula
  mass we get the number of moles.
• number of moles given mass .

gram formula mass
n 36.0g
This equation is found in table T.
18.0g
15

• ex. 1. How many moles of NaCl are in a 15.7g
  sample of NaCl?
  

given mass
number of moles
gram formula mass
x
moles mass gfm
15.7g
15.7g
x
58.5g/mole
58.5g/mole
x 0.268 moles
ex. 2. What is the mass of 1.50 moles of sulfuric
acid (H2SO4)?
m
n
gfm
x
n m gfm
1.50 moles
1.50 moles
98.1 g/mole
x
x 1.50 98.1
98.1 g/mole
x 147 g
16
H. Percent Composition the percent by mass of
each element in a compound.
mass of part

100

• composition

mass of whole

• Finding percent composition given
• experimentally obtained masses.

ex. From experimental data a compound is found to
contain 1.10g Ca, 0.880g S and 1.76g O. What is
its percent composition?
1.10g
Ca
29.4Ca
x100
3.74g
0.880g
S
23.5S
x100
3.74g
1.76g
47.1O
x100
O
3.74g
17
2. Finding Percent Composition from a Formula

• Set up Table for formula mass and divide each
  individual elements mass by the FM.
• ex. Find the percent composition by mass
  H2SO4.
• H 2 x 1.0u 2.0u
• S 1 x 32.1u 32.1u
• O 4 x 16.0u 64.0u
• 
  98.1u

2.0u
H
x 100
2.04
98.1u
32.1u
S
x 100
32.7
98.1u
64.0u
x 100
O
65.2
98.1u
18
3. Calculating Percent of Water in a Hydrate
Water in Mass of Water

x 100
a hydrate
Mass of Hydrate
Mass of Water Mass of Mass of
Hydrate
Anhydrous Salt

• Anhydrous means the salt without water, after the
  hydrate has been heated and dried.

ex. A sample of Epson Salts (MgSO4 7H2O) weighs
5.00g. After heating, the anhydrous salt
weighed 2.44g. What is the of water in the
hydrate?
19
I. Calculating Molecular Formula, given an
Empirical Formula and the Gram Molecular Mass

• The molecular formula is a multiple of the
  em-pirical formula. We need to find this
  multiple.
• ex. A compound has the empirical formula CH2
  and a gram molecular mass of 42.0g. What is its
  molecular formula?
• Trial and error method try multiples of CH2
  until you find one with the correct GMM.
• try C2H4 C 2 x 12.0g 24.0g
• H 4 x 1.0g 4.0g
• 
  28.0g too small!

20

• try C3H6 C 3 x 12.0g 36.0g
• H 6 x 1.0g 6.0g
• 
  42.0g C3H6 is right!
• 2) Calculate the Multiple Method divide the
  gram molecular mass by the mass of the empirical
  formula to get the multiple.
• Find the empirical mass C1 x 12.0g 12.0g
• 
  H2 x 1.0g 2.0g
• 
  14.0g
• Divide into molecular mass 42.0g / 14.0g 3
• Multiply the subscripts by this number
• C1x3H2x3 C3H6

21
J. Avogadros Hypothesis

• Avogadro noticed that when compounds were
  decomposed to gases, the volumes of the gases
  were always in a whole number ratio.
• ex. When water decomposes, we get twice the
  volume of Hydrogen gas as Oxygen gas.
• This ratio would seem to be the empirical formula
  of the compound if equal volumes of gases, at the
  same temperature and pressure, contained an equal
  number of molecules!

22

• This is called Avogadros Hypothesis and allowed
  Avogadro to obtain empirical formulas of many
  compounds long before molar masses were known.
• Regents example
• At the same temperature and pressure, which
  sample contains the same number of moles of
  particles as 1 liter of O2(g)?
• (1) 1 L Ne(g) (3) 0.5 L SO2(g)
• (2) 2 L N2(g) (4) 1 L H2O(l)