1. Compounds & Moles
Unit 5

2. Overview Naming The Mole Calculations Ionic Covalent Acids
Simple Organic
The Mole
Molar Mass
Mole Conversions
Calculations
Percent Composition
Empirical Formula
Molecular Formula

3. Why do we name compounds?
Think of some common compounds that you know of
H2O = water NaCl = table salt
CaCO3 = limestone
Imagine if we had to memorize common names for the millions of known compounds that we had today
…IMPOSSIBLE!
Standard system was created to name compounds
IUPAC (International Union of Pure and Applied Chemistry)

4. Indicates 18 hydrogen atoms
Chemical Formulas
Indicate the relative numbers of atoms of each kind in a chemical compound
C8H18
Indicates 8
carbon atoms
Indicates 18 hydrogen atoms

5. Molecular vs. Structural Formulas
Molecular Formula
Lists elements in a compound and how many of each element you have
Example: C2H6O
Structural Formula
Shows how atoms are “connected” in the structure
Example CH3CH2OH or CH3OCH3

6. Monatomic Ions Ions formed from a single atom Naming cations
Simply give the element’s name
Example
Ca+2 = calcium ion
Na+1 = sodium ion
Naming anions
Drop the ending of the element’s name and add “-ide”
F-1 = fluoride ion
O-2 = oxide ion

7. Binary Ionic Compounds
Ionic compound composed of 2 elements
Writing Names
Name the cation 1st
Name the anion 2nd
Example:
NaCl = sodium chloride
MgF2 = magnesium fluoride
Sr3N2 = strontium nitride

8. Binary Ionic Compounds
Writing Formulas
Example: aluminum oxide
Write the symbols for the ions side by side (cation first)
Al+3 O-2
Criss-cross the charges (use absolute value)
Al2 O3
Simplify (divide both numbers by largest common factor)
Al2O3

9. Binary Ionic Compounds
More examples (name to formula)
Calcium nitride = Ca3N2
Potassium sulfide = K2S
Magnesium oxide = MgO

10. Polyatomic Ions Electrically charged group of two or more atoms
Oxyanion – polyatomic anion that contains oxygen
General naming rules
Most common oxyanion ends in “-ate”
Example
ClO3-1 = chlorate
NO3-1 = nitrate
SO4-2 = sulfate

11. Polyatomic Ions
The number of oxygen atoms may be altered giving new endings and prefixes to oxyanions
1 more oxygen = per_______ate
Common form = _______ate
1 less oxygen = _______ite
2 less oxygens = hypo_______ite
Example
ClO4-1 = perchlorate
ClO3-1 = chlorate
ClO2-1 = chlorite
ClO-1 = hypochlorite
Notice that the charge of the oxyanion does not change (only the number of oxygen atoms)

12. Polyatomic Ions Ionic compounds (contain “ions”) Writing Name
If ion comes first, name the polyatomic ion then name the anion
If the ion comes second, name the cation then name the polyatomic ion (do not change ending)
Examples
NH4Cl = ammonium chloride
CaSO4 = calcium sulfate
Ba3(PO4)2 = barium phosphate

13. Polyatomic Ions Writing Formula
Follow same rules as binary ionic compound, but when charges are criss-crossed, use parenthesis to indicate number belongs to entire polyatomic ion
Example: calcium nitrate
Ca NO = Ca(NO3)2

14. Stock System (Ionic Compounds)
For elements that form two or more cations with different charges (example Pb+2 and Pb+4)
Uses roman numeral to indicate ion’s charge
Transition metals, Sn, and Pb use this system
Writing Formulas
Roman numeral indicates charge of the cation (use that to criss cross)
Examples
Copper (II) bromide = CuBr2
Iron (III) sulfide = Fe2S3
Tin (IV) phosphate = Sn3(PO4)4

15. Stock System (Ionic Compounds)
Writing Names
Use the anion (known charge) that the cation is bonded to and solve for the charge of the cation
Total positive charge (from cation) must equal total negative charge (from anion)
Example: VF6
Fluorine has a charge of -1
There are six fluorines bonded to the vanadium
6 × -1 = -6 so the charge of vanadium is 6
Name = vanadium (VI) fluoride

16. Stock System (Ionic Compounds)
Example 2: Sn3N2
The charge of nitrogen is -3
There are 2 nitrogen atoms
2 × -3 = -6
There are 3 tin atoms that add up to a charge of +6
+6 ÷ 3 = -2 so each tin atom has a charge of +2
Name = tin (II) nitride
Exception: some transition metals only have one charge (nickel, silver, zinc, etc.) so the roman numeral is omitted

17. Prefixes
Number
Prefix 1
mono- 2
di- 3
tri- 4
tetra- 5
penta- 6
hexa- 7
hepta- 8
octa- 9
nona- 10
deca-
Used in naming covalent compounds Indicate how many of each atom you have

18. Binary Covalent Compounds
Writing Names
Name the cation followed by the anion (-ide ending)
Use prefixes to indicate how many of each atom you have
Examples:
P4Br10 = tetraphosphorous decabromide
Si2O5 = disilicon pentoxide
Note
If an o or a are doubled, drop the o or a of the prefix
Never use mono- on cation (only on anion)

19. Binary Covalent Compounds
Writing Formulas
Prefix indicates how many of each atom you have
Do not criss-cross numbers
Examples:
Trinitrogen octachloride = N3Cl8
Arsenic tetrabromide = AsBr4

20. Summary When writing names of formulas… YES NO YES NO Is it ionic?
Is the cation a transition metal, Sn, or Pb?
Use Roman numerals
Name cation then anion (write it like it is)
Use prefixes (covalent)
When writing names of formulas…
YES NO
YES NO

21. Acids
Binary acid – contains two elements (one usually hydrogen and the other usually a halogen)
Oxyacid – acids that contain hydrogen, oxygen, and a third element (usually a nonmetal)
Usually hydrogen and a polyatomic ion

22. Acids Naming binary acids Use form of hydro_____ic acid Examples:
HF = hydrofluoric acid
HCl = hydrochloric acid

23. Acids
As the number of oxygen atoms changes in oxyacids, so does the name (just like the oxyanions)
1 more oxygen = per_______ic acid
Common form = _______ic acid
1 less oxygen = _______ous acid
2 less oxygens = hypo_______ous acid
Example
HClO4 = perchloric acid
HClO3 = chloric acid
HClO2 = chlorous acid
HClO = hypochlorous acid

24. Carbon Basis for all life.
Study of carbon compounds is called organic chemistry.
Can form single, double and triple bonds.
Long carbon chains can be produced.
Will bond with many other elements.
A HUGE number of compounds is possible (organic compounds)

25. Naming Simple Organic Compounds
Organic compounds containing only carbon and hydrogen are called hydrocarbons
Alkane – all carbons form single bonds
Alkene – carbons form double bonds
Alkyne – carbons form triple bonds
Whether a compound is an alkane, alkene, or alkyne determines the suffix (ending) in the name of the hydrocarbon

26. Naming Simple Organic Compounds
Number of carbons
determines prefix
used in name
Prefix Carbons
Meth- 1
Eth- 2
Prop- 3
But- 4
Pent- 5
Hex- 6
Hept- 7
Oct- 8
Non- 9
Dec

27. Naming Simple Organic Compounds
Examples
CH4 = methane
C2H6 = ethane
propane
propene
propyne

28. The Mole
The amount of a substance that contains as many particles as there are atoms in exactly 12 g of 12C
SI unit of amount of a substance
Abbreviated “mol”
Counting unit just like a “dozen”
1 dozen donuts is the same amount as 1 dozen books
1 mole of hydrogen atoms is the same amount as 1 mole of sodium atoms

29. Avogadro’s Number 6.022×1023 of anything is a mole
Named after Italian scientist Amadeo Avogadro
Experimentally determined number of atoms in 12 grams of 12C
How big is 602,200,000,000,000,000,000,000?
One mole of donut holes would cover the Earth 5 miles deep in the donut holes
One mole of pennies stacked on top of each other would reach from the Earth to the moon 7 times
If you started counting when you were born and never stopped until the day you died, you would never come close to reaching 6.022×1023

30. Avogadro’s Number How can that be?!
1 Liter of water contains 55.5 moles of H2O
A 5 lb bag of sugar contains 6.6 moles of sugar
How can that be?!
Atoms and molecules are so tiny that when we use units of moles (6.022×1023) it puts the particles into measurable quantities

31. Molar Mass 1 mole of hydrogen atoms = 1 mole of sodium atoms BUT…
1 mole of hydrogen atoms DOES NOT have the same mass as 1 mole of sodium atoms
Individual atoms have different masses
They are the same amount but not the same mass

32. Molar Mass The periodic table tells us the mass of 1 mole of any atom
It’s the same as the average atomic mass/relative atomic mass (decimal number on the table)
Molar Mass – mass of 1 mole of an atom or compound
Units are “grams/mole” or “g/mol”

33. Molar Mass
To find the molar mass of a compound, add the molar masses of all atoms in a compound
Also called formula mass or molecular mass (compounds only)
Example: CO2 (1 atom of C and 2 atoms of O)
1 atom C x =
2 atoms O x =
Molar mass = g/mol

34. Mole Relationships Mole Grams Atoms Molecules
6.02 x 1023
Molar Mass
To go between units of grams, moles and atoms (or molecules) use conversions!
6.022×1023 is how many atoms or molecules are in 1 mole of any substance
The molar mass is how many grams are in one mole of any substance

35. Mole Conversions How many grams are in 5.0 moles of calcium?
5.0 mole × = g
How many atoms are in 2.1 moles of xenon?
2.1 moles × = 1.26×1024 atoms g
1 mole
6.022×1023 atoms
1 mole

36. Mole Conversions
There is no way to go straight from grams to atoms or molecules in one step
Must use moles as the intermediate step
How many atoms are in 9.8 g of Pb?
9.8g × × = 2.8×1022 atoms
1 mol
207.2 g
6.022×1023 atoms
1 mole

37. Mole Conversions
When a conversion includes a compound, it will use the word molecules when a conversion includes an element, it will use the word atoms
There are still as many molecules in a mole as there are atoms
How many grams are in 3.4×1022 molecules of H2O?
First solve for molar mass of H2O
(H2O molar mass = 18.02g/mol)
3.4×1022 molecules × × = 1.0 g
1 mole
6.022×1023 molecules
18.02 g
1 mole

38. Percent Composition Percentage by mass of each element in a compound
Example: What is the percent composition of BaSO4?
Ba = 1 × = (137.3/233.4) ×100= 58.8% Ba
S = 1 × 32.1 = (32.1/233.4) ×100= % S
O = 4 × 16.0 = (64.0/233.4) ×100= % O
233.4
Multiply
by 100
Molar Mass
part ÷ total
Total molar mass

39. Divide by smallest number
Empirical Formula
Smallest whole-number ratio formula of a compound
Simplest formula
What is the empirical formula of a compound that is 27.0% sodium, 16.5% nitrogen, and 56.5% oxygen by mass?
Assume that you have a 100 gram sample
Na 27.0/22.99 = /1.17 = 1
N / = /1.17 = 1
O / = /1.17 = 3
Divide by smallest number
Molar Mass
Empirical
Formula
= NaNO3

40. Empirical Formula = Ca3N2
When numbers are too far to round, you may need to multiply all values by the same factor to make all numbers whole
What is the empirical formula of a compound that contains 40.6g of calcium and 9.5g of nitrogen?
Ca 40.6/40.1 = /0.69 = × 2 = 3
N / = /0.69 = × 2 = 2
too far to round
double both numbers to get whole numbers
Empirical
Formula
= Ca3N2

41. CH4 C3H12 Molecular Formula
Indicates actual number of atoms of each element in a compound
Multiple of empirical formula
An empirical formula can be the molecular formula, but the molecular formula is not always the empirical formula
Empirical Formula
Molecular Formula
CH4
C3H12

42. Molecular Formula
If the molecular mass is known, you can solve for the molecular formula
The molar mass of a compound with empirical formula of CH2O is g/mol. What is the molecular formula of this compound?
Molar mass CH2O = 30.02g/mol = 6
Molecular Formula = CH2O × 6 = C6H12O6
180.12
30.02