1. Unit 0 Nomenclature: Naming Chemicals
PO43-
phosphate ion
HC2H3O2
Acetic Acid
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C2H3O2-
acetate ion

2. The Period Table identifies all elements by their atomic number, which can be read off the period table.

4. History of the Periodic Table
In 1869, Dmitri Mendeleev arranged the elements according to atomic mass in an attempt to classify them.
He left spaces for elements that he predicted existed, yet hadn’t been discovered yet. This set him apart from others, and he was later proven correct.

5. Henry Moseley
In 1913, Henry Moseley developed the concept of atomic numbers. He determined the frequencies of X-rays emitted as different elements were bombarded with high energy electrons.
Each element produces X-rays of unique frequency, and the frequency increased as atomic mass increased.

6. Atomic Number
Moseley arranged these frequencies in order by assigning a unique whole number, which he called the atomic number.
He correctly identified the atomic number as the number of protons in the nucleus of the atom.
This clarified some problems that Mendeleev had; Ar and K, Te and I.

7. Periodic Table Arrangement
Periods: Horizontal Rows (7)
Indicates the energy level where the valence electrons are located.
Example: An element in Period 4 has its valence electrons in the 4th energy level.

8. Periodic Table Arrangement
Groups/Families: Modern way of labeling vertical columns; numbered from 1-18
Representative Elements(main group): groups 1,2,13-18
Transition Elements: groups 3-12 located in the center

9. Periodic Table Arrangement
Groups/Families: Antiquated way of labeling vertical columns
Roman numeral with letter A or B
Representative Elements(main group):
Group A elements
Transition Elements:
Group B elements

10. Group/Family Arrangement
Representative Elements(main group):
groups 1,2,13-18
Form ions with only 1 charge
Transition Elements: groups 3-12 located in the center
Form ions with multiple charges

11. Groups 1 and 2 Group 1 – Alkali Metals Group 2 – Alkali Earth Metals
All solids
1 valence electron/form 1+ ions
EXTREMELY reactive
Reacts with water to form H2 gas and a base
Group 2 – Alkali Earth Metals
2 valence electrons/form 2+ ions
Reactive, but not as much as Group 1

12. Group 17 – Halogens 7 valence electrons
Most reactive of the non-metals!
Gains electrons easily and forms (-) ions
Reactivity increases down the group

13. Group 18 – Noble Gases Rare or inert gases 8 valence electrons
nonreactive gases (not found in compounds in nature)
do not form ions

14. Groups 3 – 11 Transition Elements
Metallic characteristics
Form ions with multiple charges
Colored compounds

15. Broad Categories of the Periodic Table

16. Metals The elements on the left side of the table. Malleable/ Ductile
Lustrous
Great Conductors of heat and electricity
Form cations because they lose electrons
Mostly solids, Hg is liquid

17. Metals Conductors of heat and electricity
Make cations (lose e- to become + charged)
Malleable (made into sheets)
Ductile (made into wire)

18. Non-Metals The elements on the right side of the table
Brittle as solids, with no luster
Poor conductors of heat and electricity
Form anions because they gain electrons
They can be gases, liquids or solids

19. Nonmetals Are a brittle solid or a gas
Make anions (gain e- to become - charged)
Covalently bond to each
other

20. Semi-Metals (Metalloids)
Some properties are metallic, some are non-metallic and some are in-between
They are found adjacent to the darkened stair case line with the exception of Al

21. Semi-metals (AKA Metalloids)
Characteristics of both metals and nonmetals
More metallic as you go down PT

22. Chemical Formulas
A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance
An empirical formula shows the simplest
whole-number ratio of the atoms in a substance
H2O
molecular
empirical
H2O
C6H12O6
CH2O O3 O
N2H4
NH2

23. A diatomic molecule contains only two atoms
Covalent Bonding
A molecule is two or more nonmetal atoms in a definite arrangement held together by chemical bonds that share electrons H2
H2O
NH3
CH4
A diatomic molecule contains only two atoms
H2, N2, O2, Br2, HCl, CO
A polyatomic molecule contains more than two atoms
O3, H2O, NH3, CH4

24. One way to remember these elements is:
Diatomic Molecules
These seven elements occur naturally as molecules containing two atoms.
One way to remember these elements is:
Professor BrINClHOF
© 2009, Prentice-Hall, Inc.

25. The ionic compound NaCl
Ionic Bonding
A formula unit shows the smallest whole number ratio of ions in an ionic compound
Ionic Bonding- metals and nonmetals are held together by the opposite charges- electrostatic attractions.
Ionic solids are called salts .
The ionic compound NaCl

26. Naming Ions and Ionic Compounds
Formulas for ionic compounds are generally empirical formulas.
Ionic compounds are electronically neutral.

27. Ions: Atoms or groups of atoms with a charge.
Cations- positive ions - get by losing electron(s). Na1+
Anions- negative ions - get by gaining electron(s). Cl1-
Monatomic ions-single atoms with a charge
Al3+
Polyatomic ions – groups of atoms with a single charge ClO31-

28. Common Ions of Main Group Elements+1 +3
+/-4 +2 -3 -2 -1
Variable, always +

29. Look at group B elements:
Most have more than 1 charge-do not need to memorize
Silver, zinc and Cadmium only have 1 charge: must
memorize them
2.5

30. Polyatomic Ions Groups of covalently bonded atoms that have a charge.
* NO3- :nitrate ion
NO2- :nitrite ion
MUST memorize their formula and charge-see module O
NH41+ and Hg2 2+ are 2 positive polyatomics

31. Patterns for Polyatomic Ions
-ate ion
chlorate = ClO3-
-ate ion plus 1 O  same charge, per- prefix
perchlorate = ClO4-
-ate ion minus 1 O  same charge, -ite suffix
chlorite = ClO2-
-ate ion minus 2 O  same charge, hypo- prefix, -ite suffix
hypochlorite = ClO- 14

32. Writing Formulas for Ionic Compounds from Names
Formulas of ionic compounds are determined from the charges on the ions
Na S  Na S2-  Na2S
Sodium ion + sulfide ion sodium sulfide
Charge balance: = 0
Remember that all ionic compounds have no net charge, and that the charges are not written in!

33. Ionic Compounds are NEUTRAL!
Formulas are written to make the compound have a neutral charge overall.
You do NOT write the charges in the formula because they MUST cross out to accurately represent the compound.
Ex: NaF2 is INCORRECT for sodium fluoride because Na has an oxidation state of +1, and F of -1. There is a one to one ratio of Na+ to F- to make a neutral ionic compound.
Correct answer is NaF

34. So what if the oxidation numbers aren’t even?
If the oxidation numbers or charges do not balance, you can write the number of ions of each until you get the same number of each charge in total.
Yes. The “Criss Cross” method. You take the charge number from the cation, and you make it the number of anions (subscript)and take the charge on the anion, and you make that many cations (subscript).
Is there an easier way?

35. Criss-Crossing in action:
Example: Lead (II) nitride
Pb2+ N3- *the charges do not balance
Pb2+ N Pb3N2
The 2 and the 3 are brought down to the opposite element, so that there are now 3 Pb2+ ions and 2 N3- ions
This means there were 6e- transferred from the lead atoms to the nitrogen atoms; the compound is neutral

36. Writing Formulas for Ternary Ionic Compounds
Write the cation first, then the anion.
Overall charge must equal zero.
If charges cancel, just write symbols.
If not, use subscripts to balance charges.
Use parentheses to show more than one of a particular polyatomic ion.
Use Roman numerals indicate the ion’s charge when needed (transition metals)
Remember that the final formula should not have charges written in.

37. Writing Formulas, cont’d
Example:
Cr2+ PO43- *the charges do not balance
Cr2+ PO Cr3(PO4)2
The polyatomic ions is in parentheses whenever a subscript is added.
This is so that we know to count a number of those groups!

38. Examples: Write the following chemical formulas
Aluminum oxide
Calcium bromide
Sodium carbonate
Al2O3
CaBr2
Na2CO3

39. Formula of Ionic Compounds
2 x +3 = +6
3 x -2 = -6
Al2O3
Al3+
O2-
1 x +2 = +2
2 x -1 = -2
CaBr2
Ca2+
Br-
1 x +2 = +2
1 x -2 = -2
Na2CO3
Na+
CO32-

40. Is there a metal first? Naming Things:
If there is only one element present, name it. Atomic substances do not require “special” naming.
For anything with more than one element, remember that there is ONE MAIN THING to look for:
Is there a metal first?

41. So…some general help for naming:
Look to see if there is a metal first in the formula
Again, semimetals are not a classification in naming; you need to treat the elements that are on the right of the line as non-metals, and those on the left as metals.
If there are only metals, name both metals
(metallic bonding; nothing else need be done)
If ONLY the first element is a metal, then the compound is an ionic compound
Nonmetals only signify a covalent compound
There is a flow chart in your notes to help!

42. Naming Ionic Compounds:
Remember that those are compounds that have a metal first in the formula*, and then a nonmetal or a polyatomic ion.
We can handle these as simple types
binary (2 elements) compounds
Ternary(more than 2 element) compounds

43. Naming Ionic Compounds
Write the name of the cation.
Use Roman Numerals for transition metals (ONLY) after the metal
If the cation is polyatomic- name it.
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.
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44. If the Metal is a Transition Metal…
Transition metals are Type II Cations, and are elements that can have more than one possible charge. They MUST have a Roman Numeral to indicate the charge on the individual ion. 1+ or or 3+ Cu+, Cu2+ Fe2+, Fe3+ copper(I) ion iron(II) ion copper (II) ion iron(III) ion

45. FeCl3 (Fe3+) iron (III) chloride
Type II Cations
These elements REQUIRE Roman Numerals because they can have more than one possible charge:
Transition metals and the metals in groups 4A and 5A (except Ag, Zn, Cd ) require a Roman Numeral.
FeCl3 (Fe3+) iron (III) chloride
CuCl (Cu+ ) copper (I) chloride
SnF (Sn4+) tin (IV) fluoride
PbCl (Pb2+) lead (II) chloride
Fe2S (Fe3+) iron (III) sulfide

46. Type II Cations
Some Type II cations have a name using the “old” system as well as the “new system”. The old system, still widely used, adds to the root or stem of the Latin name of the metal the suffixes –ous and –ic. These represent the lower and higher charges respectively.

47. Examples: Name the following ionic compounds
BaCl2
KNO3
K2O
FeCl3
Mg(OH)2
SnO
Barium chloride
Potassium nitrate
Potassium oxide
Iron (III) chloride
Magnesium hydroxide
Tin (II) oxide

48. Examples of Older Names of Cations formed from Transition Metals (you do not have to memorize these)

49. 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

50. 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

51. Naming Molecular Compounds
Greek prefixes are used to denote the number of atoms of each element present.

52. Naming Molecular Compounds
The prefix mono- is generally omitted for the first element.
For ease of pronunciation, we usually eliminate the last letter of a prefix that ends in “o” or “a” when naming an oxide.
Example: N2O5 is dinitrogen pentoxide not dinitrogen pentaoxide

53. 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.

54. 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.

55. 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.

56. Examples: Name the following covalent compoundsHI
NF3
SO2
N2Cl4
PCl5
Hydrogen monoiodide (hydrogen iodide)
Nitrogen trifluoride
Sulfur dioxide
Dinitrogen tetrachloride
Phosphorus pentachloride

57. Rules for Writing Formulas for Molecular Compounds
Using the first part of the name, translate the name of the element into a chemical symbol and then translate the prefix if there is one to its subscript.
Using the second part of the name, translate the name of the element into a chemical symbol and then translate the prefix if there is one to its subscript.
Dichlorine trioxide Cl2O3

58. You NEVER criss-cross charges with covalent compounds.
Since you are sharing electrons, rather than giving them away/ picking them up, the charges are not relevant.

59. Acid Nomenclature Acids
Compounds that form H+ when dissolved in water.
Formulas usually begin with ‘H’.
In order to be an acid instead of a gaseous covalent compound, it must be aqueous
Meaning dissolved in water; symbolized by (aq)
Two types:
Binary
Oxyacids

60. Acid Nomenclature Flowchart

61. Naming Binary Acids To name: 1) remove the –gen ending from hydrogen
change the –ide ending on the second element to –ic.
Name the following acids
HCl(aq) hydrogen chloride → hydrochloric acid
H2S(aq) hydrogen sulfide → hydrosulfuric acid

62. Naming acids: Non-Oxy acids
If the acid doesn’t have oxygen
add the prefix “hydro-”
change the suffix “–ide” to “-ic acid”
HCl Hydrochloric acid
H2S Hydrosulfic acid
HCN Hydrocyanic acid

63. Naming Ternary 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!
HNO3
H2SO4
H2CO3
H3PO3
nitric acid
sulfuric acid
carbonic acid
phosphorous acid

64. Acid Nomenclature HBr (aq) Binary Acid H2CO3 (aq) Oxyacid
2 elements, -ide
H2CO3 (aq) Oxyacid
3 elements, -ate
H2CO2 (aq) Oxyacid
3 elements, -ite
 hydrobromic acid
 carbonic acid
 carbonous acid

66. Writing Formulas for acids
Backwards from names.
If it has hydro- in the name it has no oxygen: Binary Acid
Anion ends in “-ide”
No hydro, : Ternary Acid
anion ends in “-ate or –ite”
Write anion and add enough H’s to balance the charges.

67. A base can be defined as a substance that forms
hydroxide ions (OH-) when dissolved in water.
Most bases are metal hydroxides or NH3 (Ammonia)
There are also organic bases.
NaOH
sodium hydroxide
KOH
potassium hydroxide
Ba(OH)2
barium hydroxide

68. Name the following bases
Remember, these are just ternary ionic compounds!
Sr(OH)2
Fe(OH)3
Strontium hydroxide
Iron(III) hydroxide

69. BaI2 P4S3 Ca(OH)2 CoCO3 Na2Cr2O7 I2O5 Zn(ClO4)2 CS2 B2Cl4
Mixed Practice: Decide whether each compound is an Ionic Compound, Molecular Compound, Acid, or Base and write its name
BaI2
P4S3
Ca(OH)2
CoCO3
Na2Cr2O7
I2O5
Zn(ClO4)2
CS2
B2Cl4

70. Chromium (III) sulfate Ferric sulfite Calcium oxide Barium carbonate
Mixed Practice: Decide whether each compound is an Ionic Compound, Molecular Compound, Acid, or Base and write its formula
Dinitrogen monoxide
Potassium sulfide
Copper (II) nitrate
Dichlorine heptoxide
Chromium (III) sulfate
Ferric sulfite
Calcium oxide
Barium carbonate
Iodine monochloride

71. Organic Compounds Compounds containing carbon
Can contain elements such as hydrogen,
nitrogen, sulfur, phosphorus, oxygen &
the halogens
Have relatively low melting points,
low boiling points poor conductivity,
good solubility in other organic solvents

72. A chemical formula denotes the composition of the substance.
A molecular formula shows the exact number of atoms of each element in a molecule. C4H10
A structural formula shows not only the elemental composition, but also the general arrangements.
A condensed structural formula is a shorthand representation that leaves the bond lines out but indicates what is bonded to each carbon or other atoms

73. Hydrocarbons Hydrocarbons contain only carbon and hydrogen.
Prefixes used in naming:
1- meth 2-eth 3-prop 4-but pent hex hept
8- oct non dec
The simplest hydrocarbons are called alkanes. (CnH2n+2)
Saturated hydrocarbon chain-connected by ONLY single bonds
Tetrahedral geometry-sp3 hybridization
Least reactive of hydrocarbons
Can be in continuous chain or rings
Can have substituent groups attached to chain

74. Organic Compounds

75. Organic Compounds

76. How to Name Straight Chain Alkanes
Count the number of carbons in the chain.
Use the prefix stem to represent the number of carbons
Use the suffix –ane
CH4
CH3CH2CH3

77. How to Name Branched Chain Alkanes
Identify & name the longest, continuous carbon chain. This is called the parent chain which is named using a root to indicate the # of carbons followed by “ane”.
Identify & name the substituent groups located off the continuous chain:
Alkyl groups are named using a root to indicate the # of carbons in the group followed by “yl”. Examples -CH3 is methyl -CH2CH3 is ethyl CH2CH2CH3 is propyl Halogen groups are name like this F: fluoro -Cl: chloro
(The prefixes di-, tri-,etc are used to indicate multiple identical substituents.)
3. The position of the substituent groups are specified by numbering the longest chain starting at the end closest to the branching (the substituent group).

78. How to Name Branched Chain Alkanes
4. Designate the location of each substituent group by an appropriate number & name. Identify where in the chain the substituent group is located with a number placed before the name.
If two are on the same carbon use that number twice. Separate the number and the name with a hyphen and commas to separate numbers from each other.
5. The location & name of the substituent groups are followed by the root alkane name. The substituents are listed in alpha order. Prefixes are not considered when alphabetizing.

79. Examples
CH3CH(CH3)CH2CH3

80. Alkenes Hydrocarbons can also be alkenes. (CnH2n)
Unsaturated hydrocarbon chain-carbons are connected by AT LEAST 1 double bond
Trigonal Planar geometry-sp2 hybridization
Can have substituent groups attached to chain

81. How to Name Alkenes
1. Indicate where the double bond is by numbering the longest chain that contains both the carbons of the double bond starting with the one closest to the bond. The smaller number of the two carbon atoms of the double bond is used as the double bond locator.
2. Then use the prefix stem to indicate the number of carbon atoms in the chain & the suffix “ene” ending.
3. Follow rules from “alkanes” to name substituent groups attached to alkenes.
If the double bond is in the center of the chain, the nearest substituent group rule is used to determine where numbering begins
If more than 1 double bond is present, the compound is named diene, triene, etc. Each double bond is given a locator number.

82. Alkene Examples
C3H6

83. Alkynes Hydrocarbons can also be alkynes. (CnH2n-2)
Unsaturated hydrocarbon chain-carbons are connected by AT LEAST 1 triple bond
Linear geometry-sp hybridization
Can have substituent groups attached to chain

84. How to Name Alkynes
1. Indicate where the triple bond is by numbering the longest chain that contains both the carbons of the triple bond starting with the one closest to the bond. The smaller number of the two carbon atoms of the triple bond is used as the triple bond locator.
2. Then use the prefix stem to indicate the number of carbon atoms in the chain & the suffix “ene” ending.
3. Follow rules from “alkanes” to name substituent groups attached to alkenes.
If the triple bond is in the center of the chain, the nearest substituent group rule is used to determine where numbering begins
If more than 1 triple bond is present, the compound is named diene, triene, etc. Each triple bond is given a locator number.

85. Alkyne Examples
C2H2

86. Cyclo Alkanes
All carbons are single bonded to two other carbons in a ring structure.
To name them:
Rings are numbered to give the smallest
substituent numbers possible.
2. Largest substituent groups are given the
lowest locator number

87. Aromatic Hydrocarbons
Nonpolar hydrocarbons with six membered carbon rings with delocalized electrons
The simplest molecule: Benzene: C6H6: six carbon ring with alternating double bonds

88. CycloAlkane Examples

89. Functional Groups
Many organic compounds contain groups of atoms known as functional groups, which often determine a molecule’s reactivity.
You are responsible for recognizing ALCOHOLS & CARBOXYLIC ACUDS

90. Alcohols
The –OH (hydroxyl group) is covalent bonded to C so it does not ionize to produce hydroxide ions
To name:
Indicate the number of carbon atoms to
which the –OH group is attached.
2. Drop the “e” from the name of the alkane,
and add the suffix –ol
3. Start counting in the chain closest to the
–OH group

91. Alcohols

92. Carboxylic Acids Occur in plants and animals
The COOH group must always be on the end and its “C” is always the #1 C in the chain
To name:
drop the “e” from the name of the alkane, and add the suffix –oic followed by the word “acid”.

93. Carboxylic Acids

94. Name the following Examples
C4H9OH
CH3CH2CH2CH2COOH

95. Structural Isomers Two or more compounds with the same chemical
formula but a different arrangement of atoms in
the molecule and thus different properties.
There are 2 isomers of butane
CH3CH2CH2CH and CH3CH(CH3)CH3
(butane) (2-methylpropane)
What are the 5 isomers of hexane?