1. Chapter 3
Chemical Compounds

2. 3.1 Types of Chemical Compounds and Their Formulas

3. Molecular Compounds
Molecular compounds are composed of molecules and contain only nonmetals.
Electrons are shared
AKA Covalent Compounds
Represented by chemical formulas

4. The ratio of the masses of carbon and hydrogen, C:H in methane is
1. 4:1
2. 1:4
3. 3:1
4. 1:3
5. 1:1
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5. The ratio of the masses of carbon and hydrogen, C:H in methane is
1. 4:1
2. 1:4
3. 3:1
4. 1:3
5. 1:1
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6. The ratio of the masses of carbon and hydrogen, C:H in ethane is
1. 4:1
2. 1:4
3. 3:1
4. 1:3
5. 1:1
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7. The ratio of the masses of carbon and hydrogen, C:H in ethane is
1. 4:1
2. 1:4
3. 3:1
4. 1:3
5. 1:1
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8. Ions When atoms lose or gain electrons, they become ions.
Cations are positive and are formed by elements on the left side of the periodic chart.
Anions are negative and are formed by elements on the right side of the periodic chart.

9. Figure: UN
Title:
Ionization of sodium.
Caption:
A sodium atom loses an electron to form a sodium ion.

10. Figure: UN
Title:
Formation of a chloride ion.
Caption:
A chlorine atom gains an electron to form a chloride ion.

11. Ionic Bonds
Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

12. A mole of solid sodium chloride, salt, contains
22.99 g of sodium and g of chlorine A B
x1023 g of sodium and 6.02x1023 g of chloride
u of sodium and u of chloride
g of sodium and g of chloride.
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13. A mole of solid sodium chloride, salt, contains
22.99 g of sodium and g of chlorine A B
x1023 g of sodium and 6.02x1023 g of chloride
u of sodium and u of chloride
g of sodium and g of chloride.
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14. A molecule of solid sodium chloride, salt, contains
22.99 g of Na and g of Cl A B
u of Na+ and u of Cl-
g of Na and g of Cl
4. None of the above
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15. A molecule of solid sodium chloride, salt, contains
22.99 g of sodium and g of chlorine A B
u of sodium and u of chloride
g of sodium and g of chloride.
4. None of the above; sodium chloride is an ionic compound, as such, it does not exist in discrete molecular units as is characteristic of covalent compounds.
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16. Would you expect the following to be ionic or molecular:
N2O
Na2O
CaCl2
SF4
CBr4
FeS
P4O6
PbF2

17. Types of Formulas
Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound.
Molecular formulas give the exact number of atoms of each element in a compound.
Structural Formulas show the order in which atoms are bonded

18. Types of Formulas
Structural formulas show the order in which atoms are bonded.
Perspective drawings also show the three-dimensional array of atoms in a compound.

19. Types of Formulas
Formula Unit: the smallest neutral collection of ions (most reduced, for ionic)

20. For the ball and stick model of naphthalene to the right, the empirical formula is
1. C10H8
2. C4H5
3. C5H4
4. CH
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21. For the ball and stick model of naphthalene to the right, the empirical formula is
1. C10H8
2. C4H5
3. C5H4
4. CH
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22. For the ball and stick model of pyrimidine to the right, the molecular formula is:
1. C4N2H4
2. C2N2H2
3. C2NH2
4. (CH)4N2
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23. For the ball and stick model of pyrimidine to the right, the molecular formula is:
1. C4N2H4
2. C2N2H2
3. C2NH2
4. (CH)4N2
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24. 3.2 The Mole Concept and Chemical Compounds

25. Formula Mass vs. Molar Mass vs. Molecular Mass
Formula Mass: the mass of a formula unit (amu)
Molecular Mass: the mass of a molecule (amu)
Molar Mass: the mass of one mole of a compound (grams)
H2O: Molecule Mass amu (1 molecule)
Molar Mass grams (1 mole)

26. Molecular Formulas
Diatomic molecules: H2 O2 N2 F2 Cl2 Br2 I2 Molecules: P4 (White Phosphorus) S8 (Sulfur) Distinguish between molecule and atom!

27. Example 3-1A
How many grams of MgCl2 would you need to obtain 5.0 x Cl- ions?

28. Example 3-2A
Gold has a density of g/cm3. A piece of gold foil is 2.50 cm on each side and 0.100mm thick. How many atoms of gold are in this piece of gold foil?

29. Example 3-3A
Halothane: C2HBrClF3 How many grams of Br are contained in mL of halothane (d=1.871g/mL)

30. 3.3 Composition of Chemical Compounds

31. Percent Composition
One can find the percentage of the mass of a compound that comes from each of the elements in the compound by using this equation:
% element =
(number of atoms)(MM of element)
(MM of the compound)
x 100

32. Percent Composition So the percentage of carbon in ethane is…
(2)(12.0 g/mol)
(30.0 g/mol)
24.0 g/mol
30.0 g/mol =
x 100
= 80.0%

33. Example 3-4A
What are the mass percent composition in C10H16N5P3O13?

34. Calculating Empirical Formulas
One can calculate the empirical formula from the percent composition

35. Establishing Formulas From % Comp
Assume 100g, %  grams
Grams  Moles
Divide by smallest moles
Get whole numbers in most reduced form
Ration= MM Molecular/ MM Empirical
Multiply Ratio by Empirical to get Molecular Formula

36. Example 3-5A
Sorbitol is a sweetener that has a molecular mass of 182 u and percent composition of 39.56% C 7.74% H % O What are the empirical and molecular formulas?

37. Calculating Empirical Formulas
The compound para-aminobenzoic acid (you may have seen it listed as PABA on your bottle of sunscreen) is composed of carbon (61.31%), hydrogen (5.14%), nitrogen (10.21%), and oxygen (23.33%). Find the empirical formula of PABA.

38. Calculating Empirical Formulas
Assuming g of para-aminobenzoic acid,
C: g x = mol C
H: g x = 5.09 mol H
N: g x = mol N
O: g x = mol O
1 mol
12.01 g
14.01 g
1.01 g
16.00 g

39. Calculating Empirical Formulas
Calculate the mole ratio by dividing by the smallest number of moles:
C: =  7
H: =  7
N: = 1.000
O: =  2
5.105 mol
mol
5.09 mol
1.458 mol

40. Calculating Empirical Formulas
These are the subscripts for the empirical formula:
C7H7NO2

41. Combustion Analysis
Compounds containing C, H and O are routinely analyzed through combustion in a chamber like this
C is determined from the mass of CO2 produced
H is determined from the mass of H2O produced
O is determined by difference after the C and H have been determined

42. Example
The combustion of 5.00 grams of an alcohol produces 9.55 g of CO2 and 5.87 g of H2O. Find the empirical formula.
Solution: First you need to find the individual masses of the elements.
The general equation would look like this:
CxHyOz + O2  CO2 + H2O

43. The mass of the carbon in the CO2 all came from the alcohol so…
9.55 gCO2 X g C g C
44.01 gCO2
The mass of the hydrogen in H2O all came from the alcohol so…
5.87 g H2O X g H g H
18.02 g H2O

44. Because the mass of the oxygen in both CO2 and H2O is derived from both the alcohol and the oxygen from combustion you need to find the mass of oxygen only in the alcohol….
By subtracting the mass of the carbon and hydrogen from the total mass of the compound.
5.00 g – (2.61 g g) = 1.73 g O

45. Now you can determine the number of moles of each element in the formula…..
Moles of C = 2.61 g X 1.00 mole C mol C
12.01 g C
Moles of H = g X 1.00 mole H mol H
1.01 g H
Moles of O = 1.73 g O X 1.00 mole O mol O
16.00 g O

46. From the number of moles we find the mole ratios,….
C = 0.217/0.108 = 2.01
H = 0.651/ = 6.03
O = 0.108/ = 1.00
Thus the formula will be C2H6O or C2H5OH

47. Practice Problem
1.540 g of an organic acid burns completely to produce g CO2 and g H2O. Find the empirical formula. If the molecular mass is 90.0 grams what is the molecular formula?
C gCO2 X 12.01/44.01 = g
H g H2O X 2.02/18.02 = g
O –( ) = g
Empirical formula= CH2O
Molecular formula = (CH2O)x
X= molecular mass/empirical mass
X= 90/30= C3H6O3
C /12.01=
H / 1.01 =
O /16.00=

48. Example 3-6B
Combustion of a g sample of thiophene, a carbon-hydrogen-sulfur compound yields g CO2, g H2O and g SO2 as the only combustion products. What is the empirical formula?

49. 3.4 Oxidation States
Tell the number of electrons gained or lost when forming compounds
Handout

50. Common Oxidation States

51. Rules for Assigning Oxidation States (OS)
The OS of a free element is 0
The total OS of all atoms in a compound is 0
The total OS of all atoms in a ion is equal to the charge of the ion
Group 1 = +1, Group 2 = +2
Fluorine = -1
Hydrogen usually +1
Oxygen usually -2
In Binary Ionic : Group 17 = -1, Group 16 = -2, Group 15 = -3

52. Exceptions Rule 6: H bonded to metals; LiH, NaH, CaH2
Rule 7: O-F Bonds OF2, H2O2, and KO2

53. Example
What are the oxidation states of: S8 Cr2O72- MgCl2 Al2O3 MnO4-

54. The oxidation number for N in nitric acid, HNO3, is
1. 1
2. 2
3. 3
4. 4
5. 5
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55. The oxidation number for N in nitric acid, HNO3, is
1. 1
2. 2
3. 3
4. 4
5. 5
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56. 3.5 Naming Compounds: Organic and Inorganic Compounds

57. Cations 1. Cations from metal atoms have the same name as the metal
Na+ sodium ion Ba2+ barium ion
2. Use Roman numerals to indicate the charge on a metal ion that can form more than one positive charge
Fe2+ Iron (II) ion Fe3+ Iron(III) ion

58. Cations
Latin names for some metals still are used and the ending –ous and –ic are used to indicate the charge
Fe2+ ferrous ion Fe3+ ferric
Sn2+ stannous ion Sn4+ stannic ion
Other latin root names include:
Plumbous
Auric
Cuprous

59. Cations Cations from nonmetal atoms end in-ium
NH4+ ammonium ion H3O+ hydronium ion

60. Anions Monatomic anions are named are named by adding –ide at the end
H- Hydride S2- sulfide P3- phosphide
Some polyatomics have –ide endings
OH- hydroxide
CN- cyanide
O22- peroxide

61. Anions Polyatomic anions containing oxygen end in –ite or -ate
The –ate ending is used for the most common oxyanion of the element . The –ite ending is used for the oxyanion that has the same charge but one O atom less.
NO3- nitrate NO2- nitrite
SO42- sulfate SO32- sulfite

62. Anions
Prefixes are used when there is a series of four oxyanions. (usually the halogens)
Per- is used to indicate one more O than the –ate ending and hypo- is used for one less O than the –ite ending.
ClO4- perchlorate
ClO3- chlorate
ClO2- chlorite
ClO- hypochlorite

63. Anions
Anions with H+ are named by adding the prefix hydrogen or dihydrogen
HCO3- hydrogen carbonate (bicarbonate)
HSO4- dihydrogen sulfate (bisulfate)

64. Writing Formulas
Because compounds are electrically neutral, one can determine the formula of a compound this way:
The charge on the cation becomes the subscript on the anion.
The charge on the anion becomes the subscript on the cation.
If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor.

65. Potassium dichromate is used in breathalyzers
Potassium dichromate is used in breathalyzers. When it comes in contact with alcohol vapor it turns from orange to green. It is an ionic compound where the polyatomic anion has the formula Cr2O72-. What is the chemical formula for potassium dichromate?
1. KCr2O7
2. K(Cr2O7)2
3. K2Cr2O7
4. K2(Cr2O7)3
5. K3(Cr2O7)2
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66. Potassium dichromate is used in breathalyzers
Potassium dichromate is used in breathalyzers. When it comes in contact with alcohol vapor it turns from orange to green. It is an ionic compound where the polyatomic anion has the formula Cr2O72-. What is the chemical formula for potassium dichromate?
1. KCr2O7
2. K(Cr2O7)2
3. K2Cr2O7
4. K2(Cr2O7)3
5. K3(Cr2O7)2
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67. Sodium phosphate is an active component in some constipation medicines and enemas. The chemical formula for sodium phosphate is
1. NaPO4
2. Na2PO3
3. Na2PO4
4. Na3PO4
5. Na3(PO4)2
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68. Sodium phosphate is an active component in some constipation medicines and enemas. The chemical formula for sodium phosphate is
1. NaPO4
2. Na2PO3
3. Na2PO4
4. Na3PO4
5. Na3(PO4)2
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69. Marble consists of primarily of calcium carbonate
Marble consists of primarily of calcium carbonate. Acids react with and dissolve marble as evident by this marble statue eroded by acid rain. The chemical formula for calcium carbonate is
1. CaCO3
2. Ca2CO3
3. CaCO4
4. Ca3(CO3)2
5. Ca(CO3)2
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70. Marble consists of primarily of calcium carbonate
Marble consists of primarily of calcium carbonate. Acids react with and dissolve marble as evident by this marble statue eroded by acid rain. The chemical formula for calcium carbonate is
1. CaCO3
2. Ca2CO3
3. CaCO4
4. Ca3(CO3)2
5. Ca(CO3)2
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71. Common Cations

72. Common Anions

73. Inorganic Nomenclature
Write the name of the cation.
If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion.
If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.

74. Name the following compounds: K2SO4 Ba(OH)2 FeCl3
Write the chemical formulas for the following compounds:
potassium sulfide,
calcium hydrogen carbonate,
nickel(II) perchlorate.

75. Acid Nomenclature
If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- :
HCl: hydrochloric acid
HBr: hydrobromic acid
HI: hydroiodic acid

76. Acid Nomenclature
If the anion in the acid ends in -ite, change the ending to -ous acid:
HClO: hypochlorous acid
HClO2: chlorous acid

77. Acid Nomenclature
If the anion in the acid ends in -ate, change the ending to -ic acid:
HClO3: chloric acid
HClO4: perchloric acid

78. Nomenclature of Binary Covalent Compounds
The less electronegative atom is usually listed first.
A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however.)

79. Nomenclature of Binary Compounds
The ending on the more electronegative element is changed to -ide.
CO2: carbon dioxide
CCl4: carbon tetrachloride

80. Nomenclature of Binary Compounds
If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one:
N2O5: dinitrogen pentoxide

81. 3.6 Names and Formulas of Inorganic Compounds

82. Organic compounds Alkanes-Carbon atoms bonded to four other atoms
Methane CH4
Ethane C2H6
Propane C3H8
Butane C4H10
Alkanes with 5 or more carbons use the following prefixes
Penta
Hexa
Hepta
Octa
Nona
Deca

83. Organic functional groups
An alcohol is an alkane with an –OH group and is named by adding an –ol ending
Methanol
In naming the carbon with the functional group is identified by a number
2-Propanol