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COMPOUNDS & MOLES
Unit 5
Overview

Naming

 Ionic
 Percent
 Covalent
 Acids
 Simple

Organic
The Mole
 Molar
Calculations
Mass
 Mole Conversions
Composition
 Empirical Formula
 Molecular Formula
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)
Chemical Formulas

Indicate the relative numbers of atoms of each kind
in a chemical compound
Indicates 8
carbon atoms
C8H18
Indicates 18
hydrogen atoms
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
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”
Example


F-1 = fluoride ion
O-2 = oxide ion
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
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
Binary Ionic Compounds

More examples (name to formula)
 Calcium
nitride = Ca3N2
 Potassium
sulfide = K2S
 Magnesium
oxide = MgO
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
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)
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
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:
Ca+2
calcium nitrate
NO3-1
=
Ca(NO3)2
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
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

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
Prefixes
Used in naming
covalent compounds
Indicate how many of
each atom you have
Number Prefix
1
mono-
2
di-
3
tri-
4
tetra-
5
penta-
6
hexa-
7
hepta-
8
octa-
9
nona-
10
deca-
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)

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
Summary
When writing names
of formulas…
Is it ionic?
YES
NO
Is the cation a
transition metal,
Sn, or Pb?
YES
Use Roman
numerals
NO
Name cation
then anion (write
it like it is)
Use prefixes
(covalent)
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
Acids

Naming binary acids
 Use
form of
hydro_____ic acid
 Examples:
 HF
= hydrofluoric acid
 HCl = hydrochloric acid
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
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)
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
Naming Simple Organic Compounds
Prefix
MethEthPropButPentHexHeptOctNonDec-
Carbons
1
2
3
4
5
6
7
8
9
10
Number of carbons
determines prefix
used in name
Naming Simple Organic Compounds

Examples
 CH4
= methane
 C2H6 = ethane
propane
propene
propyne
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
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
Avogadro’s Number


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
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
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
 Units
Mass – mass of 1 mole of an atom or compound
are “grams/mole” or “g/mol”
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 12.011
2 atoms O x 15.9994
Molar mass
=
=
=
12.011
31.9988
44.010 g/mol
Mole Relationships
6.02 x 1023
Atoms
Molecules

Molar Mass
Mole
Grams
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

Mole Conversions

How many grams are in 5.0 moles of calcium?
40.078 g
5.0 mole ×
= 200.39 g
1 mole

How many atoms are in 2.1 moles of xenon?
6.022×1023 atoms
2.1 moles ×
= 1.26×1024 atoms
1 mole
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?
1 mol
6.022×1023 atoms
9.8g ×
×
= 2.8×1022 atoms
207.2 g
1 mole
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
18.02 g
1 mole
molecules × 6.022×1023 molecules × 1 mole = 1.0 g
Percent Composition

Percentage by mass of each element in a compound
Example: What is the percent composition of BaSO4?
Molar Mass
part ÷ total
Multiply
by 100
Ba = 1 × 137.3 = 137.3 (137.3/233.4) ×100= 58.8% Ba
S=
1 × 32.1 = 32.1 (32.1/233.4) ×100= 13.8% S
O = 4 × 16.0 = 64.0 (64.0/233.4) ×100= 27.4% O
233.4
Total molar mass
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
Molar Mass
Divide by
smallest number
Na 27.0/22.99 = 1.17 /1.17 = 1
N 16.5/14.01 = 1.18 /1.17 = 1
O 56.5/16.00 = 3.53 /1.17 = 3
Empirical
= NaNO3
Formula
Empirical Formula

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?
too far to round
Ca 40.6/40.1 = 1.01 /0.69 = 1.5 × 2 = 3
N 9.5/14.01 = 0.69 /0.69 = 1 × 2 = 2
Empirical
Formula = Ca3N2
double both numbers
to get whole numbers
Molecular Formula

Indicates actual number of atoms of each element
in a compound
 Multiple

of empirical formula
Empirical
Formula
Molecular
Formula
CH4
C3H12
An empirical formula can be the molecular formula, but the
molecular formula is not always the empirical formula
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 180.12 g/mol. What is the molecular
formula of this compound?
Molar mass CH2O = 30.02g/mol
180.12
30.02 = 6
Molecular Formula = CH2O × 6 = C6H12O6