Chapter 7

Chemical Formulas and Chemical Compounds

Section 1 – Chemical Names and Formulas

Significance of Chemical Formula

!Chemical formula shows number and types of atoms in compound

!Ex. C6H12O6

!Carbon = 6 atoms

!Hydrogen = 12 atoms

!Oxygen = 6 atoms

!Ionic compound made of mixture of + and – ions held by attraction

!Combination of cation (+) and anion (-)

Al2(SO4)3

SO4 x 3

Aluminum = 2

Sulfur = 3

Oxygen = 12

Monoatomic Ions

!Monoatomic ions – ions formed from single atom

!Group 1 – lose one e-, form +1

!Group 2 – lose two e-, form +2

!Groups 15, 16, 17 gain electrons to form anions

!-3, -2, and -1, respectively

!Not all main-group elements form ions easily

!C and Si form covalent bonds

!Transition metals can form +1, +2, +3, +4

Naming Monoatomic Ions

!Cations named same as element name

!K+ = potassium

!Mg2+ = magnesium

!Anions named by dropping ending and adding –ide

!F- = fluoride

!N3- = nitride

Binary Ionic Compounds

!Binary compounds ! compounds made of two different elements !Total positive and negative charge must be equal Mg2+ Br- +2 + -1 = +1

+2 + 2(-1) = 0

Formula = MgBr2

Naming Binary Ionic Compounds

!Nomenclature ! naming system !Combine names of cations and anions !Cation (+) name ALWAYS comes first !Then anion (-) name

Al2O

3

Aluminum oxide

Practice Problem 1

!Write formulas for the binary ionic compounds formed between the following elements:

!a. potassium and iodine

! KI !b. magnesium and chlorine

! MgCl2

!c. sodium and sulfur

! Na2S

!d. aluminum and sulfur

! Al2S3

!e. aluminum and nitrogen

! AlN

Practice Problem 2

!Name the binary ionic compounds:

!a. AgCl

! Silver chloride

!b. ZnO

! Zinc oxide

!c. CaBr2

! Calcium bromide

!d. SrF2

! Strontium fluoride

!e. BaO

! Barium oxide

!f. CaCl2

! Calcium chloride

Stock System

!Some elements form two or more cations with different charges

!Stock system uses Roman numerals to show ion’s charge

!Roman numeral included in () right after metal name

Fe2+ Iron (II)

CuCl2

Copper (II) chloride

Name of cation

Roman numeral indicating charge of cation

Name of anion

Practice Problem 1

!Write the formula and give the name for the compounds formed between the following ions:

!a. Cu2+ and Br−

! CuBr2

! copper(II) bromide

!b. Fe2+ and O2−

! FeO

! iron(II) oxide

!c. Pb2+ and Cl−

! PbCl2

! lead(II) chloride

!d. Hg2+ and S2−

! HgS

! mercury(II) sulfide

!e. Sn2+ and F−

! SnF2

! tin(II) fluoride

!f. Fe3+ and O2−

! Fe2O3

! iron(III) oxide

Polyatomic Ions

!Polyatomic ion ! ion containing more than one atom

!Most are oxyanions ! polyatomic ions that contain oxygen

NO2- NO3

-

!Ion with more oxygen ends in –ate (nitrate) !Ion with less oxygen ends in –ite (nitrite)

Hypo- and Hyper-

!One oxyanion family has 4 members

ClO- ClO2

- ClO3

- ClO4

- Hypochlorite Chlorite Chlorate Perchlorate

!Naming same as binary ionic compounds !Cation first, anion second

Practice Problem 1

! Write formulas for the following ionic compounds:

! a. sodium iodide

! NaI

! b. calcium chloride

! CaCl2

! c. potassium sulfide

! K2S

! d. lithium nitrate

! LiNO3

!e. copper(II) sulfate

! CuSO4

!f. sodium carbonate

! Na2CO

3

!g. calcium nitrite

! Ca(NO2)2

!h. potassium perchlorate

! KClO4

Practice Problem 2

!Give the names for the following compounds:

!a. Ag2O

! silver oxide

!b. Ca(OH)2

! calcium hydroxide

!c. KClO3

! potassium chlorate

!d. NH4OH

! ammonium hydroxide

!e. FeCrO4

! iron(II) chromate

!f. KClO

! potassium hypochlorite

Naming Binary Molecular Compounds

!Molecular compounds made of covalently bonded molecules

!Usually happens between 2 NON-metals

Prefix System

!Uses prefixes to show how many atoms presentNumber Prefix

1 Mono-

2 Di-

3 Tri-

4 Tetra-

5 Penta-

6 Hexa-

7 Hepta-

8 Octa-

9 Nona-

10 Deca-

Examples

!CCl4 ! carbon tetrachloride

!P4O10 ! tetraphosphorous decoxide

!There are rules for determining which ion comes first

Rules

1. The less-electronegative element is given first. If only 1 atom, no prefix used

2. Second element named by combining (a) prefix, (b) element root name, (c) –ide

3. The o or a at the end of prefix dropped when word following beings with vowel

Practice Problem

!Name the following binary molecular compounds:

!a. SO3

! sulfur trioxide

!b. ICl3

! iodine trichloride

!c. PBr5

! phorphorous pentabromide

Acids

!Most acids either binary acids or oxyacids

!Binary acids ! acids that are made of 2 elements

!Usually H and one of the halogens

!Oxyacids ! acids that contain H, O, and a 3rd element (usually nonmetal)

Naming Acids

!Binary acids – hydro(element root)ic acid

!Ex. HCl – hydrochloric acid

!Oxyacids

!– with less oxygen, (element root)ous acid !– with more oxygen, (elementroot)ic acid

!Ex. HNO2 = nitrous acid

! HNO3 = nitric acid

Formula Name

HF Hydrofluoric acid

HCl Hydrochloric acid

HBr Hydrobromic acid

HI Hydroiodic acid

H3PO4 Phosphoric acid

HNO2 Nitrous acid

HNO3 Nitric acid

H2SO3 Sulfurous acid

H2SO4 Sulfuric acid

CH3COOH Acetic acid

HClO Hypochlorous acid

HClO2 Chlorous acid

HClO3 Chloric acid

HClO4 Perchloric acid

H2CO3 Carbonic acid

Salts

!Salt ! an ionic compound made of a cation and an anion from an acid

!Table salt – NaCl – contains anion from HCl

!Some salts have anions where one or more H atoms from acid are kept

!Named by adding hydrogen OR prefix bi- to anion name

HCO3- hydrogen carbonate ion

bicarbonate ion

Practice Problem 1

!Write formulas for the compounds formed between the following:

!a. aluminum and bromine

! AlBr3

!b. sodium and oxygen

! Na2O

!c. magnesium and iodine

! MgI2

!d. Pb2+ and O2−

! PbO

!e. Sn2+ and I−

! SnI2

!f. Fe3+ and S2−

! Fe2S

3

!g. Cu2+ and NO3−

! Cu(NO3)2

!h. NH4+ and SO42−

! (NH4)2SO

4

Practice Problem 2

!Name the following compounds using the Stock system:

!a. NaI

! sodium iodide

!b. MgS

! magnesium sulfide

!c. CaO

! calcium oxide

! d. K2S

! potassium sulfide

!e. CuBr

! copper (I) bromide

! f. FeCl2

! iron (II) chloride

Practice Problem 3

!Write formulas for each of the following compounds: !a. barium sulfide ! BaS !b. sodium hydroxide ! NaOH !c. lead(II) nitrate

! Pb(NO3)2

!d. potassium permanganate

! KMnO4

!e. iron(II) sulfate

! FeSO4

!f. diphosphorus trioxide

! P2O

3

!g. disulfur dichloride

! S2Cl

2

!h. carbon diselenide

! CSe2

!i. acetic acid

! CH3COOH

!j. chloric acid

! HClO3

!k. sulfurous acid

! H2SO

3

! l. phosphoric acid

! H3PO

4

Section 2 – Oxidation Numbers

The charges on the ions making an ionic compound reflect the electron distribution of the compound. In order to indicate the general distribution of electrons among the bonded atoms in a molecular compound or a polyatomic ion, oxidation numbers, also called oxidation states, are assigned to the atoms composing the compound or ion.

Unlike ionic charges, oxidation numbers do not have an exact physical meaning. However, oxidation numbers are useful in naming compounds, in writing formulas, and in balancing chemical equations. And they are helpful in studying certain types of chemical reactions.

Assigning Oxidation Numbers

!Shared electrons assumed to belong to more-electronegative atom in each bond

Rules

1. Atoms in a pure element have Ox# of zero. 1. Na, O2, P4 = 0

2. More-electronegative element in binary molecular compound assigned the number equal to negative charge it has as ion. Less electronegative assigned equal to positive charge it has as ion.

3. Fluorine always is -1 because it is most electronegative.

4. Oxygen has -2 in almost all compounds, except

Peroxides (H2O2, O = -1)

Compounds with halogens (O = +)

5. H = +1 in all compounds with more electronegative element, H = -1 in compounds with metals

6. Sum of Ox# of all atoms in neutral compound = zero

7. Sum of Ox# of all atoms in polyatomic ion = charge of ion

8. Ox# of ions in ionic compounds = its charge

Sample Problem 1

!Assign Ox# to each atom in UF6.

!Start by placing known Ox#s above appropriate elements.

!From rules, we know F is always -1

!Multiply known Ox#s by appropriate number of atoms and place totals underneath matching elements.

!There are 6 F atoms, 6 x -1 = -6

!UF6 is molecular

!According to guidelines, sum of Ox#s must be zero

!Total positive Ox#s must be +6

!Divide total positive Ox#s by number of atoms

!+6 ÷ 1 = +6

H2SO4

+1 -2 H

2 S O

4

+2 -8

+1 +6 -2 H

2 S O

4

+2 +6 -8

Practice Problem

! Assign oxidation numbers to each atom in the following compounds or ions:

! a. HCl

! +1, -1

! b. CF4

! +4, -1

! c. PCl3

! +3, -1

!d. SO2

!+4, -2

!e. HNO3

!+1, +5, -2

!f. KH

!+1, -1

!g. P4O10

!+5, -2

!h. HClO3

!+1, +5, -2

!i. N2O5

!+5, -2

!j. GeCl2

!+2, -1

Stock System

!Remember there were two ways to name covalent compounds !Stock system

!Prefix system

!Some nonmetals can have more than one oxidation number

!Listed in Table A-15 in handout

Group 14 Carbon -4, +2, +4

Group 15 Nitrogen Phosphorous

-3, +3, +5 -3, +3, +5

Group 16 Sulfur -2, +4, +6

Group 17 Chlorine Bromine Iodine

-1, +1, +3, +5, +7 -1, +1, +3, +5, +7 -1, +1, +3, +5, +7

!Can use Roman numerals to show oxidation number

Prefix System Stock System

PCl3 Phosphorous trichloride

Phosphorous(III) chloride

PCl5 Phosphorous pentachloride

Phosphorous(V) chloride

N2O Dinitrogen oxide Nitrogen(I) oxide

NO Nitrogen monoxide Nitrogen(II) oxide

PbO2 Lead dioxide Lead(IV) oxide

Mo2O3 Dimolybdenum trioxide

Molybdenum(III) oxide

Practice Problem 1

! Assign oxidation numbers to each atom in the following compounds or ions:

! a. HF

! +1, -1

! b. CI4

! +4, -1

! c. H2O

! +1, -2

! d. PI3

! +3, -1

!e. CS2

!+4, -2

!f. Na2O

2

!+1, -1

!g. H2CO

3

!+1, +4, -2

! h. NO2 −

!+3, -2

!i. SO42−

!+6, -2

!j. ClO2 −

!+3, -2

!k. IO3 −

!+5, -2

Practice Problem 2

!Name each of the following binary molecular compounds according to the Stock system:

! a. CI4

! carbon(IV) iodide

!b. SO3

! sulfur(VI) oxide

! c. As2S

3

! arsenic(III) sulfide

!d. NCl3

! nitrogen(III) chloride

Section 3 – Using Chemical Formulas

As you have seen, a chemical formula indicates the elements as well as the relative number of atoms or ions of each element present in a compound. Chemical formulas also allow chemists to calculate a number of characteristic values for a given compound. In this section, you will learn how to use chemical formulas to calculate the formula mass, the molar mass, and the percentage composition by mass of a compound.

The Mole

!SI unit for an amount of substance (like 1 dozen = 12) !Avogadro’s number ! the number of particles in exactly one

mole of a pure substance !6.022 x 1023

!How big is that ? !If 5 billion people worked to count the atoms in one mole of an

element, and if each person counted continuously at a rate of one atom per second, it would take about 4 million years for all the atoms to be counted

Molar Masses

!Molar mass ! the mass of one mole of a pure substance

!Written in unit g/mol

!Found on periodic table (atomic mass)

!Ex. Molar mass of H = 1.008 g/mol Molar Mass H2O = ?

H = 2 x 1.01 g/mol O = 1 x 16.00 g/mol

H2O = 18.02 g/mol

Practice Problem

! Find the molar mass of each of the compounds

! a. Al2S

3

! 150.17 g/mol

! b. NaNO3

! 85.00 g/mol

! c. Ba(OH)2

! 171.35 g/mol

Molar Mass as Conversion Factor

!g/mol can be used as a conversion factor to change from moles to mass

!What is the mass in grams of 2.50 mol of oyxgen gas?

Practice Problem 1

!How many moles of compound are there in the following?

!a. 6.60 g (NH4)2SO4

! 0.0500 mol !b. 4.5 kg Ca(OH)2

! 61 mol

Practice Problem 2

!How many molecules are there in the following?

!a. 25.0 g H2SO4

! 1.53 × 1023 molecules

!b. 125 g of sugar, C12H22O11

! 2.20 × 1023 molecules

What is the mass in grams of 3.50 mol of the element copper, Cu?

222 g Cu

What is the mass in grams of 2.25 mol of the element iron, Fe?

126 g Fe

What is the mass in grams of 0.375 mol of the element potassium, K?

14.7 g K

What is the mass in grams of 0.0135 mol of the element sodium, Na?

0.310 g Na

What is the mass in grams of 16.3 mol of the element nickel, Ni?

957 g Ni

1. How many moles of lead, Pb, are in 1.50 × 1012 atoms of lead?

2.49 × 10−12 mol Pb

2. How many moles of tin, Sn, are in 2500 atoms of tin?

4.2 × 10−21 mol Sn

3. How many atoms of aluminum, Al, are in 2.75 mol of aluminum?

1.66 × 1024 atoms Al

1. What is the mass in grams of 7.5 × 1015 atoms of nickel, Ni?

7.3 × 10−7 g Ni

2. How many atoms of sulfur, S, are in 4.00 g of sulfur?

7.51 × 1022 atoms S

3. What mass of gold,Au, contains the same number of atoms as 9.0 g of aluminum,Al?

66 g Au

Percent Composition

!Percentage composition ! percentage by mass of each element in a compound

!Divide mass of element in sample of compound by total mass of sample, multiply by 100

Sample Problem

!Find the percentage composition of copper(I) sulfide, Cu2S.

1. Analyze

!Given:

!formula, Cu2S

!Unknown:

!percentage composition of Cu2S

2. Plan

!formula → molar mass → mass percentage of each element

!The molar mass of the compound must be found

!Then the mass of each element present in one mole of the compound is used to calculate the mass percentage of each element.

3. Compute

Molar mass:

Cu = 2 x 63.55 g Cu = 127.1 g Cu

S = 1 x 32.07 g S = 32.07 g S

159.2 g/mol Cu2S

Sample Problem 2

!As some salts crystallize from a water solution, they bind water molecules in their crystal structure. Sodium carbonate forms such a hydrate, in which 10 water molecules are present for every formula unit of sodium carbonate. Find the mass percentage of water in sodium carbonate decahydrate, Na2CO3•10H2O, which has a molar mass of 286.14 g/mol.

1. Analyze

!Given:

!chemical formula, Na2CO

3•10H

2O

!molar mass of Na2CO

3•10H

2O

!Unknown:

!mass percentage of H2O

2. Plan

!chemical formula → mass H2O per mole of Na

2CO

3•10H

2O

→ % water

!The mass of water per mole of sodium carbonate decahydrate must first be found.

!This value is then divided by the mass of one mole of Na

2CO

3•10H

2O.

3. Compute

!1 mol Na2CO

3•10H

2O contains 10 mol of H

2O

Section 4 – Determining Chemical Formulas

!When a new substance is synthesized or is discovered, it is analyzed to show its percentage composition

!From this data, the empirical formula is then determined

!An empirical formula consists of the symbols for the elements combined in a compound, with subscripts showing the smallest whole-number mole ratio of the different atoms in the compound

!For an ionic compound, the formula unit is usually the compound’s empirical formula

!For a molecular compound, however, the empirical formula does not necessarily indicate the actual numbers of atoms present in each molecule

!For example, the empirical formula of the gas diborane is BH3, but the molecular formula is B2H6

!In this case, the number of atoms given by the molecular formula corresponds to the empirical ratio multiplied by two.

Calculation of Empirical Formulas

!Determine empirical formula from percent composition

!Change percent composition to grams

!Ex. 29.9% H …. 29.9 g H

!Use molar mass to change grams to mole

!Divide all moles by smallest mole amount

!Result is the subscript for the empirical formula

Example

!Percent composition of diborane is 78.1% B and 21.9% H

!In 100.0 g sample of diborane, there is 78.1 g B and 21.9 g H

!Mass composition of each element converted to composition in moles by dividing by molar mass

!7.22 mol B to 21.7 mol H

!Not a ratio of smallest whole numbers

!Divide each number of mol by smallest number in ratio

Rounding, empirical formula is BH3

Sample Problem 1

!Quantitative analysis shows that a compound contains 32.38% sodium, 22.65% sulfur, and 44.99% oxygen. Find the empirical formula of this compound.

1. Analyze

!Given: percentage composition:

!32.38% Na,

!22.65% S, and

!44.99% O

!Unknown: empirical formula

2. Plan

percentage composition

mass

moles

smallest whole-number mole ratio of atoms

3. Compute

!Mass composition (mass of each element in 100.0 g sample):

!32.38 g Na,

!22.65 g S,

!44.99 g O

!Divide each mole by smallest mole in ratio

Na2SO4

4. Evaluate

!Calculating the percentage composition of the compound based on the empirical formula determined in the problem reveals a percentage composition of 32.37% Na, 22.58% S, and 45.05% O

!These values equal 100%

Practice Problem 1

!Analysis of a 10.150 g sample of a compound known to contain only phosphorus and oxygen indicates a phosphorus content of 4.433 g. What is the empirical formula of this compound?

!PO3

Practice Problem 2

!A compound is found to contain 63.52% iron and 36.48% sulfur. Find its empirical formula.

!FeS

Practice Problem 3

!Find the empirical formula of a compound found to contain 26.56% potassium, 35.41% chromium, and the remainder oxygen.

!K2Cr

2O

7

Practice Problem 4

!Analysis of 20.0 g of a compound containing only calcium and bromine indicates that 4.00 g of calcium are present. What is the empirical formula of the compound formed?

!CaBr2

Calculating Molecular Formulas

!Empirical formula contains smallest whole number ratio

!Molecular formula is actual formula for compound

!Empirical ! CH2

!Molecular ! C2H4 (ethene), C3H6 (cyclopropane)

!Relationship between empirical and molecular:

x(empirical formula) = molecular formula

!x is whole-number multiple !Must know molecular mass

!Ex. Experiment shows molar mass of diborane (BH3) is 27.67 g/mol !Empirical formula mass = 13.84 g/mol

Practice Problem 1

!The empirical formula of a compound of phosphorus and oxygen was found to be P2O5. Experimentation shows that the molar mass of this compound is 283.89 g/mol. What is the compound’s molecular formula?

!P4O10

Practice Problem 2

!Determine the molecular formula of the compound with an empirical formula of CH and a formula mass of 78.110 g/mol.

!C6H

6

Practice Problem 3

!A sample of a compound with a formula mass of 34.00 g/mol is found to consist of 0.44 g H and 6.92 g O. Find its molecular formula.

!H2O

2