1. Chemical Nomenclature

2. Octet Rule Atoms tend to achieve electron configuration of Noble Gases
Octet = Eight
Noble Gases have eight electrons in their highest energy level
General Equation for Noble Gases is S2P6

3. Atoms of Metallic Elements tend to lose valence electron/s, leaving an octet in the next lowest energy level
Atoms of a Non-Metallic Element tend to gain a valence electron/s to achieve an Octet
There are EXCEPTIONS to the Octet Rule

4. Diatomic Molecules
These eight elements occur naturally as molecules containing two atoms.
Astatine is considered a diatomic
© 2009, Prentice-Hall, Inc.

5. Elements that exist as Diatomic Molecules
Most elements, except for the noble gases, do not exist as single atoms.
The halogens, along with hydrogen, nitrogen, and oxygen, exist naturally as diatomic molecules when in their element form (not as compounds). H2 F2
Cl2
Br2 I2
N O2
Figure 2.22

6. Ions Atoms or groups of atoms with a charge
Cations- positive ions - get by losing electrons(s)
Anions- negative ions - get by gaining electron(s)
Ionic bonding- held together by the opposite charges
Ionic solids are called salts

7. Ions
Ions differ from atoms in that they have a charge; the number of electrons is either greater than, or less than, the number of protons.
Cations are positively charged. They have fewer electrons than in the neutral atom.
Anions are negatively charged. They have more electrons than in the neutral atom.

8. Ions and the Periodic Table
The noble gases are the most stable (least reactive) elements on the periodic table.
Their stability is associated with the number of electrons they contain.
Many atoms in the main-group elements gain or lose electrons to obtain the same number of electrons as the nearest noble gas.
Metals tend to lose electrons, and therefore become cations.
Nonmetals tend to gain electrons, thereby becoming anions.

9. Even though atoms and cations have the same name, there are many chemical differences between metals and their cations.
Example:
Na Metal; reacts explosively in water
Na Cation; quite unreactive

10. K+1 Ca+2 Cations Positive ions. Formed by losing electrons.
More protons than electrons.
Metals form cations.
K+1
Has lost one electron
Ca+2
Has lost two electrons

11. Formation of Mg2+ Cation
Figure 2.14

12. F-1 O-2 Anion A negative ion. Has gained electrons.
Non metals can gain electrons.
Charge is written as a super script on the right.
F-1
Has gained one electron
O-2
Has gained two electrons

13. Formation of N3 Anion
Figure 2.14

14. Charges on ions
For most of the Group A elements, the Periodic Table can tell what kind of ion they will form from their location.
Elements in the same group have similar properties.
Including the charge when they are ions.

15. Monatomic Ions- consist of a single atom with a positive or negative charge resulting from the loss or gain of one or more valence electrons
Groups 1a, 2a, and 3a lose electrons and form cations
Aluminum is the only common group 3a element to lose electrons and form a cation

16. Non-metals tend to gain electrons and form an anion.
Groups 5a, 6a, and 7a form anions
In group 5a, there are three non-metals which form anions
N3-, P3-, & As3-

17. Majority of elements in 4a & 0 do not form ions
Group 0 rarely forms compounds
Ordinarily, two non-metals from group 4a, C & Si are found in molecular compounds

18. +1 +2 -3 -2 -1

19. Laws Conservation of Mass
Law of Definite Proportion- compounds have a constant composition by mass.
They react in specific ratios by mass.
Multiple Proportions- When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers.

20. Law of Constant Composition Joseph Proust (1754–1826)
This is also known as the law of definite proportions.
It states that the elemental composition of a pure substance never varies.
© 2009, Prentice-Hall, Inc.

21. Law of Conservation of Mass
The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.
© 2009, Prentice-Hall, Inc.

22. Compounds Follow the Law of Definite Proportion.
Have a constant composition.
Have to add the same number of atoms every time.
Two types.

23. Molecular Compounds Molecular compounds Made of molecules.
Made by joining nonmetal atoms together into molecules.

24. Ionic Compound Ionic Compounds Made of cations and anions.
Metals and nonmetals.
The electrons lost by the cation are gained by the anion.
The cation and anions surround each other.
Smallest piece is a FORMULA UNIT.

25. Ionic Compounds
Figure 3.8

26. Formula Unit
Formula Unit- lowest whole-number ratio of the ions in the compound
Example
Na+Cl-
Ratio is 1:1
The formula unit is NaCl

27. Formula Unit
The smallest whole number ratio of atoms in an ionic compound.
Ions surround each other so you can’t say which is hooked to which.

28. Two Types of Compounds Ionic Molecular Smallest piece Formula Unit
Molecule
Types of elements
Metal and Nonmetal
Nonmetals
liquid or gas
State
solid
Melting Point
High >300ºC
Low <300ºC

29. Ionic and Molecular Compounds

30. Activity: Identifying Ionic and Molecular Compounds
Identify each compound as ionic or molecular.
CCl4
CaF2
SF6
CuCO3
H2O

31. 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.
© 2009, Prentice-Hall, Inc.

32. Writing Formulas Two sets of rules, ionic and covalent
To decide which to use, decide what the first word is.
If is a metal or polyatomic use ionic.
If it is a non-metal use covalent

33. 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.
© 2009, Prentice-Hall, Inc.

34. Ionic Formulas Charges must add up to zero
get charges from table, name of metal ion, or memorized from the list
use parenthesis to indicate multiple polyatomics

35. Ionic Formulas Sodium nitride sodium- Na is always +1
nitride - ide tells you it comes from the table
nitride is N-3

36. Na+1 N-3 Ionic Formulas Sodium nitride sodium- Na is always +1
nitride - ide tells you it comes from the table
nitride is N-3
doesn’t add up to zero
Na+1
N-3

37. Na3N Na+1 N-3 Ionic Formulas Sodium nitride sodium- Na is always +1
nitride - ide tells you it comes from the table
nitride is N-3
doesn’t add up to zero
Need 3 Na
Na3N
Na+1
N-3

38. Writing Formulas Write the formula for calcium chloride.
Calcium is Ca+2
Chloride is Cl-1
Ca+2 Cl-1 would have a +1 charge.
Need another Cl-1
Ca+2 Cl2-1

39. Crisscross
Switch the numerical value of the charges
Ba2+
N3- 2 3
Ba3 N2
Reduce ratio if possible

40. Polyatomic Ions
Polyatomic Ion- Tightly bound groups of atoms that behave as a unit and carry a charge
Unlike monatomic ions; Sulfate anion is composed of 1 Sulfur atom and 4 oxygen atoms
These five atoms form a Sulfate Anion
It has a –2 charge an is written SO42-

41. Polyatomic Ions
A polyatomic ion consists of a group of atoms with an overall net charge.
These two are also called oxoanions because they contain oxygen attached to some other element.
The most common polyatomic cation is NH4+.
Figure 3.14

42. Polyatomic anions either end in ITE or ATE
Out of the two similar polyatomic ions, the polyatomic with less Oxygens ends in ite
Example:
Sulfite and Sulfate
Sulfite; SO32-
Sulfate; SO42-

43. There are three exceptions to the Polyatomic Rule
1) Ammonium NH The only positive polyatomic ion
2) Cyanide CN Ends in IDE
3) Hydroxide OH Ends in IDE

44. Write the formulas for these
Lithium sulfide
tin (II) oxide
tin (IV) oxide
Magnesium fluoride
Copper (II) sulfate
Iron (III) phosphide
gallium nitrate
Iron (III) sulfide

45. Ionic Compounds Sodium sulfite calcium iodide Lead (II) oxide
Lead (IV) oxide
Mercury (I) sulfide
Barium chromate
Aluminum hydrogen sulfate
Cerium (IV) nitrite

46. Write the formulas for these
Ammonium chloride
ammonium sulfide
barium nitrate

47. Naming compounds Two types Ionic - metal and non metal or polyatomics
Covalent- we will just learn the rules for 2 non-metals

48. There are two methods for naming cations with multiple charges
The Stock System and Classical System
The Stock system is the preferred method

49. Stock System The stock system uses roman numerals in
( ). The ( ) indicate the numerical charge of the cation.
Example:
Fe2+ Name: Iron(II)
There is no space between the name and the parenthesis
Cu1+ Name:
Copper(1)

50. Classical System
The classical system uses the root word with different suffixes as the end of the word
OUS- is used to name the cation with the lower of the two ionic charges
IC- is used to name the cation with the higher of the two ionic charges

51. Example:
Fe2+ and Fe3+
Name: Ferrous
Name: Ferric
What is the problem with the classical system?

52. The classical system does not tell you the charge of the ion.
The name only tells you which cation is either larger or smaller out of the pair

53. Few transition metals have only one ionic charge
These three elements don’t have roman numerals next to there name
Exceptions:
Ag+
Cd2+
Zn2+

54. Ionic compounds
If the cation is monoatomic- Name the metal (cation) just write the name.
If the cation is polyatomic- name it
If the anion is monoatomic- name it but change the ending to -ide
If the anion is poly atomic- just name it
practice

55. Naming Binary Ionic Compounds
Write the name of CuO
Need the charge of Cu
O is -2
copper must be +2
Copper (II) chloride
Name CoCl3
Cl is -1 and there are three of them = -3
Co must be +3 Cobalt (III) chloride

56. Naming Binary Ionic Compounds
Write the name of Cu2S.
Since S is -2, the Cu2 must be +2, so each one is +1.
copper (I) sulfide
Fe2O3
Each O is x -2 = -6
3 Fe must = +6, so each is +2.
iron (III) oxide

57. Ternary Ionic Compounds
Will have polyatomic ions
At least three elements (3 capital letters)
Still just name the ions
NaNO3
CaSO4
CuSO3

58. Ternary Ionic Compounds
LiCN
Fe(OH)3
(NH4)2CO3
NiPO4

59. Ionic Compounds Have to know what ions they form
off table, polyatomic, or figure it out
CaS
K2S
AlPO4
K2SO4
FeS
CoI3

60. Ionic Compounds
Fe2(C2O4)
MgO
MnO
KMnO4
NH4NO3
Hg2Cl2
Cr2O3

61. Ionic Compounds
KClO4
NaClO3
YBrO2
Cr(ClO)6

62. Writing names and Formulas
Molecular Compounds
Writing names and Formulas

63. Molecules & Molecular Compounds
Elements are the building materials of the substances that make up all living and nonliving things
Only about 100 elements but there are millions of different compounds made from their atoms
Thus, naming compounds is an essential skill in chemistry

64. In nature, only Noble Gases tend to exist as isolated atoms.
They are monatomic; that is, they consist of single atoms
Many elements found in nature are in the form of molecules
Molecule- is the smallest electrically neutral unit of a substance that still has the properties of the substance
Molecules are made up of two or more atoms that act as a unit

65. Atoms of different elements may combine chemically to form compounds
In many compounds, the atoms combine to form molecules.
Molecular Compounds- Compounds composed of molecules
Molecular Compounds tend to have relativity low melting and boiling points
Many of these compounds thus exist as gases or liquids at room temperature.

66. Molecular compounds made of just nonmetals
smallest piece is a molecule
can’t be held together because of opposite charges.
can’t use charges to figure out how many of each atom

67. Naming Covalent Compounds
Two words, with prefixes
Prefixes tell you how many.
mono, di, tri, tetra, penta, hexa, septa, nona, deca
First element whole name with the appropriate prefix, except mono
Second element, -ide ending with appropriate prefix
Practice

68. Nomenclature of Binary 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) .
© 2009, Prentice-Hall, Inc.

69. Nomenclature of Binary Compounds
The ending on the more electronegative element is changed to -ide.
CO2: carbon dioxide
CCl4: carbon tetrachloride
© 2009, Prentice-Hall, Inc.

70. 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
© 2009, Prentice-Hall, Inc.

71. Naming Covalent Compounds
CCl4
N2O4
XeF6
N4O4
P2O10

72. Name These
N2O
NO2
Cl2O7
CBr4
CO2
BaCl2

73. Covalent compounds The name tells you how to write the formula duh
Sulfur dioxide
diflourine monoxide
nitrogen trichloride
diphosphorus pentoxide

74. Write formulas for these
diphosphorus pentoxide
tetraiodide nonoxide
sulfur hexaflouride
nitrogen trioxide
Carbon tetrahydride
phosphorus trifluoride
aluminum chloride

75. More Names and formulas

76. Acids Substances that produce H+ ions when dissolved in water
All acids begin with H
Two types of acids
Oxyacids
non oxyacids

77. Naming acids If the formula has oxygen in it
write the name of the anion, but change
ate to -ic acid
ite to -ous acid
Watch out for sulfuric and sulfurous
H2CrO4
HMnO4
HNO2

78. Naming acids If the acid doesn’t have oxygen add the prefix hydro-
change the suffix -ide to -ic acid
HCl
H2S
HCN

80. Formulas for acids Backwards from names
If it has hydro- in the name it has no oxygen
anion ends in -ide
No hydro, anion ends in -ate or -ite
Write anion and add enough H to balance the charges.

81. Formulas for acids hydrofluoric acid dichromic acid carbonic acid
hydrophosphoric acid
hypofluorous acid
perchloric acid
phosphorous acid

82. Hydrates Some salts trap water crystals when they form crystals
these are hydrates.
Both the name and the formula needs to indicate how many water molecules are trapped
In the name we add the word hydrate with a prefix that tells us how many water molecules

83. Hydrates
In the formula you put a dot and then write the number of molecules.
Calcium chloride dihydrate = CaCl2·2H2O
Chromium (III) nitrate hexahydrate = Cr(NO3)3· 6H2O