1. Chemical Formula & Naming
Chapter 7

2. CCl MgCl2
Guess the name of each of the above compounds based on the formulas written.
What kind of information can you discern from the formulas?
Guess which of the compounds represented is molecular(covalent) and which is ionic.
Chemical formulas form the basis of the language of chemistry and reveal much information about the substances they represent.

3. example: octane — C8H18 Carbons = 8 Hydrogens = 18
A chemical formula indicates the relative number of atoms of each kind in a chemical compound.
example: octane — C8H18
Carbons = 8
Hydrogens = 18

4. example: aluminum sulfate — Al2(SO4)3
Parentheses surround the polyatomic ion to identify it as a unit. The subscript 3 refers to the unit.
Note also that there is no subscript for sulfur: when there is no subscript next to an atom, the subscript is understood to be 1.

5. Monatomic Ions
Many main-group elements can lose or gain electrons to form ions.
Ions formed from a single atom are known as monatomic ions.
example: To gain a noble-gas electron configuration, nitrogen gains three electrons to form N3– ions.
Some main-group elements tend to form covalent bonds instead of forming ions.
examples: carbon and silicon

6. Naming Monatomic Ions
Monatomic cations are identified by the element’s name.
examples:
K+ is called the potassium cation
Mg2+ is called the magnesium cation
For monatomic anions, the ending of the element’s name is dropped, and the ending -ide is added to the root name.
F– is called the fluoride anion
N3– is called the nitride anion

9. Binary compounds
Compounds composed of two elements are binary compounds.
In a binary ionic compound:
positive charges = negative charges
The formula for a binary ionic compound can be written cation 1st , anion 2nd .
example: magnesium bromide
Ions combined: Mg2+, Br–, Br–
Chemical formula: MgBr2

10. Example
Aluminum
Chloride Al +3 Cl -1
AlCl3

11. Example
Sodium
Sulfide Na +1 S -2
Na2S

12. Example
Magnesium
Oxide Mg +2 O -2
MgO

13. Sample Problem
Write the formulas for the binary ionic compounds formed between the following elements:
a. iodine and zinc
b. zinc and sulfur

14. Polyatomic Ions
Many common polyatomic ions are oxyanions—polyatomic ions that contain oxygen.
Examples:
PO4, SO4, NO2

15. Polyatomic Ions (cont.)
Some elements can combine with oxygen to form more than one type of oxyanion.
example: nitrogen
NO3- NO2-
The name of the ion with the greater number of oxygen atoms ends in -ate. The smaller number of oxygen atoms ends in -ite.
nitrate nitrite

16. Mg SO4 MgSO4 Example Magnesium Sulfate +2 -2
-ate and -ite endings are polyatomics
Magnesium
Sulfate Mg +2
SO4 -2
MgSO4

17. Example
Aluminum
Nitrite Al +3
NO2 -1
Al(NO2)3

18. Example
Calcium
Hydroxide Ca +2 OH -1
Ca(OH)2

19. Example
Iron (III)
Oxide Fe +3 O -2
Fe2O3

20. Sample Problem
Write the formula for tin(IV) sulfate.

21. Example
Ba3N2
barium
nitride

22. Example
Na3PO4
sodium
phosphate

23. CuO Example Copper (II) Oxide Copper ( ) Oxide Copper (I) – Cu+1
Which copper is it?
-2 X 1 atom = -2
Copper (II) Oxide
Copper ( ) Oxide
Copper (I) – Cu+1
Copper (II) – Cu+2

24. FeCl3 Example Iron (III) Chloride Iron (II) – Fe+2 Iron (III) – Fe+3
Which iron is it?
-1 X 3 atom = -3
Iron (III) Chloride
Iron (II) – Fe+2
Iron (III) – Fe+3

25. Cu3PO4 Example Copper (I ) Phosphate Copper ( ) Phosphate
Which copper is it?
-3 X 1 atom = -3
+3/3 Cu atoms = +1
Copper (I ) Phosphate
Copper ( ) Phosphate
Copper (I) – Cu+1
Copper (II) – Cu+2

26. Naming Binary Molecular Compounds
Molecular compounds are composed of individual covalently bonded units.
Naming molecular compounds is based on the use of prefixes.
examples: CCl4 — carbon tetrachloride CO — carbon monoxide CO2 — carbon dioxide

27. Prefixes
1 – mono 2 – di 3 – tri 4 – tetra 5 – penta 6 – hexa 7 – hepta 8 – octa 9 – nona 10 - deca

28. Naming (cont.)
Rule – Use prefixes to tell how many of each atom; 1st atom name from periodic table, 2nd element gets –ide ending Exception – If only 1 of first atom, NO mono-

29. Prefixes

30. Example
Dichlorine
heptabromide
Cl2
Br7

31. S O3 Example Sulfur trioxide
Note: no “mono” when only one of first element

32. Sample Problem a. Give the name for As2O5.
b. Write the formula for oxygen difluoride.

33. Example
F3P4
trifluorine
tetraphosphide

34. Example
CO2
carbon
dioxide

35. Example CO
carbon
monoxide

36. Warm-UP Potassium carbonate Dinitrogen pentachloride CBr4 AuP
Mercury(I) selenide

37. Acids
Can be either:
Binary acids are acids that consist of two elements, usually hydrogen and a halogen.
Oxyacids are acids that contain hydrogen, oxygen, and a third element (usually a nonmetal).

38. Naming Acids – Create the rules
Examples of binary acids:
Hydrochloric acid HCl
Hydrofluoric acid HF
Examples of oxyacids:
phosphoric acid H3PO4
nitric acid HNO3
sulfuric acid H2SO4

39. Acid Chart ____-ide ____-ate ____-ite hydro-___-ic acid _____-ic acid
Ion ending
____-ide
____-ate
____-ite
Acid name is…
hydro-___-ic acid
_____-ic acid
_____-ous acid

40. HBr Example bromide = hydrobromic acid ____-ide ____-ate ____-ite
Ion ending
____-ide
____-ate
____-ite
Acid name is…
hydro-___-ic acid
_____-ic acid
_____-ous acid

41. H3N Example nitride = hydronitric acid ____-ide ____-ate ____-ite
Ion ending
____-ide
____-ate
____-ite
Acid name is…
hydro-___-ic acid
_____-ic acid
_____-ous acid

42. HNO3 Example nitrate = nitric acid ____-ide ____-ate ____-ite
Ion ending
____-ide
____-ate
____-ite
Acid name is…
hydro-___-ic acid
_____-ic acid
_____-ous acid

43. Hydrosulfuric acid Example H+1 S-2 = H2S ____-ide ____-ate ____-ite
Ion ending
____-ide
____-ate
____-ite
Acid name is…
hydro-___-ic acid
_____-ic acid
_____-ous acid

44. H2SO3 Example sulfite = sulfurous acid ____-ide ____-ate ____-ite
Ion ending
____-ide
____-ate
____-ite
Acid name is…
hydro-___-ic acid
_____-ic acid
_____-ous acid

45. All Mixed Up Determine compound type
Ionic – starts with metal or NH4
Covalent – starts with nonmetal
Acid – starts with hydrogen
Use proper rules to write formula/name

46. Section 7-3 The chemical formula for water is H2O.
How many atoms of hydrogen and oxygen are there in one water molecule?
How might you calculate the mass of a water molecule, given the atomic masses of hydrogen and oxygen?

47. Chemical formulas allow chemists to calculate a number of characteristic values for a compound:
formula mass
molar mass
percentage composition

48. The formula mass of any molecule, formula unit, or ion is the sum of the average atomic masses of all atoms represented in its formula.
example: formula mass of water, H2O
average atomic mass of H: 1.01 amu
average atomic mass of O: amu
average mass of H2O molecule: amu

49. formula mass of KClO3 = 122.55 amu
Sample Problem:
Find the formula mass of potassium chlorate, KClO3
formula mass of KClO3 = amu

50. A compound’s molar mass is numerically equal to its formula mass.
Ex.) the molar mass of pure calcium, Ca, is g/mol because one mole of calcium atoms has a mass of g.
Ex.) molar mass of H2O molecule: g/mol

51. molar mass of Ba(NO3)2 = 261.35 g/mol
Sample Problem G
What is the molar mass of barium nitrate, Ba(NO3)2?
molar mass of Ba(NO3)2 = g/mol

52. Molar Mass as a conversion factor

53. Sample Problem
What is the mass in grams of 2.50 mol of oxygen gas (O2)?
moles O2  grams O2
amount of O2 (mol)  molar mass of O2 (g/mol) = mass of O2 (g)

54. Sample Problem Ibuprofen, C13H18O2, is the active ingredient in many
nonprescription pain relievers. Its molar mass is
g/mol.
If the tablets in a bottle contain a total of 33 g of ibuprofen, how many moles of ibuprofen are in the bottle?
How many molecules of ibuprofen are in the bottle?
What is the total mass in grams of carbon in 33 g
of ibuprofen?

55. Percent Composition
To find the mass percentage of an element in a compound, the following equation can be used.
Mass of element/mass of total sample x 100 = percent comp.

56. Sample Problem
Find the percentage composition of copper(I) sulfide, Cu2S.

57. More problems
From data: What is the percent composition of a compound made from g Na and 77.4 g O?
From formula: Find the percent composition of sodium sulfate.

58. Empirical Formula The simplest ratio of atoms in a compound
Need percent composition to find it

59. Empirical formula, cont.
Example: Find the empirical formula of a compound that is 79.9% C and 20.1 % H.
Assume 100 g
Convert to moles
Divide by smallest number of moles
*If all are close to whole numbers, stop
*If NOT, multiply all to make them all whole
4. Write the formula

60. Empirical formula, cont.
Ex: Find the empirical formula of a compound that is 17.6% Na, 39.7% Cr, and 42.7% O.
Assume 100 g
Convert to moles
Divide by smallest number of moles
*If all are close to whole numbers, stop
*If NOT, multiply all to make them all whole
4. Write the formula

61. Example: diborane
The percentage composition is 78.1% B and 21.9% H. What is the empirical formula?

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

63. Molecular Formulas NOT the simplest ratio
Is a MULTIPLE of the empirical formula
Ex.
C2H4  CH2
Molecular formula
Empirical
formula

64. Molecular formulas, cont.
Ex. A compound that is 58.8% C, 9.8% H, and 31.4% O, has a molecular/formula mass of 204 g/mol. Find its formula mass.
Find the empirical formula
Find mass of the empirical formula
Divide molecular mass by empirical mass to find multiplier
Write the molecular formula

65. Chapter Review
Pg. 251
#’s 2-8, 10-12, 14, 23, 30-32, 35-38

66. Mole review 1 mole = 6.02E23 atoms 1 mole = ____g
1 mole = 6.02E23 molecules
1 mole = 6.02E23 formula units
How many formula units are in 4.3 grams silver nitrate?
How many moles is 3.43 X 1024 molecules carbon dioxide?
How many grams is 2.65 X 1023 atoms of aluminum?