Download Chemical Composition Notes

CHEMICAL FORMULAS & COMPOSITION
John Dalton said compounds form by the combination
of atoms in the ratio of small whole numbers.
He devised a symbol for each known
element/compound to help in writing chemical
equation
____________ - smallest unit of a compound that
retains the chemical characteristics of the compound;
characteristics of the constituent elements are lost.
____________________ - uses symbols and subscripts to
represent the composition of the molecule.
Strictest sense—covalently bonded
EX:
________________ – shows how the atoms are grouped
and identifies important parts of the molecule
EX:
____________________ – shows how all the atoms are
attached within the molecule.
Significance of a Chemical Formula
Al2(SO4)3
IONS AND IONIC COMPOUNDS
_______ - formed when electrons are lost or gained in
ordinary chemical reactions; affect size of atoms dramatically
__________ - (+) ions; often metals since metals lose
electrons to become positively charged
________ - (—) ions; often nonmetals since nonmetals gain
electrons to become negatively charged
For the metals 1A through 3A, for 1B and 2B, and for the
metals of group 4A, it is equal to the group number
Transition metals 3B through 8B, figure a 2+ or 3+ charge (use
roman numerals)
Maximum number of e- gained by a nonmetal is equal to 8
minus group number of the element.
Hydrogen can either gain or lose one electron, depending on
the other elements it encounters
Noble gases do not lose or gain electrons, except in rare
cases!!
ELEMENTS THAT EXIST AS MOLECULES
Pure hydrogen, nitrogen, oxygen and the halogens exist as
DIATOMIC molecules under normal conditions. MEMORIZE!!!
Br2 I2 N2 Cl2 H2 O2 F2
P4 S8 –
Carbon - diamond and graphite 
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_________________ are the polyatomic ions containing oxygen.
ends in –ate and –ite
prefix hypo – if it contains one O less than –ite
prefix per – if it contains one O more than –ate
PRACTICE PROBLEMS
Using these guidelines, predict possible charges for ions formed from
a)K
d)V
b)Se
e)Co
c)Be
f)Cs
OXIDATION NUMBERS
__________________ – a number assigned to an atom in a
molecular compound or ion that indicates the general distribution of
electrons among the bonded atoms
Rules in Assigning Oxidation Numbers
Oxidation Number of Pure Element =
Binary Molecular Compound
Less electronegative atom =
More electronegative atom =
Fluorine =
Compounds formed from ions:
Cation’s symbol is written first (on the left)
Ionic or not Ionic?
Ionic if one of the elements is a metal.
Ionic if a metal is combined with a nonmetal.
Ionic if a metal is combined with polyatomic ion
PRACTICE PROBLEM
Oxygen =
Exceptions:
Peroxides (w/ 1A, 2A metals, H2O2) =
with Fluorine (Ex. OF2), =
Hydrogen
Chromium is a transition metal and so can form ions with at least two different
charges. Write the formulas of the compounds formed between chromium and
sulfur.
NAMES OF CATIONS
Naming Cations: Usually Metals
Monatomic Metal Cation — simply the name of the metal from which
it is derived.
EX: Al3+ is the aluminum ion
Transition metals form more than one ion;
Roman Numerals inside parenthesis follow the ion’s name.
EX: Cu2+ is copper (II)
Mercury (I) is an exception  it is Hg22+ — two Hg+ bonded
together.
NH4+ is ammonium
As Hydrides (with metals – NaH) =
Algebraic sum of the oxidation numbers of all atoms in a neutral
compound =
EX: KMnO4
Algebraic sum of the oxidation numbers of all atoms in a neutral
compound =
EX: KMnO4
Algebraic sum of the oxidation numbers of all atoms in a
polyatomic ion =
EX: (SO4)2Although rules 1 through 7 apply to covalently bonded atoms,
oxidation numbers can also be assigned to atoms in ionic
compounds
NAMES OF ANIONS
________________ - add the suffix —ide to the stem of the
nonmetal’s name. Halogens are called the halides.
_________________ – group of covalently bonded atoms that carry a
net charge; units of atoms behaving as one entity
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PRACTICE PROBLEMS
Assign Oxidation Numbers to each atom in the following compounds:
(IONIC) Metal (or NH4+) + Nonmetal or Polyatomic
ion
HCl
FIXED CHARGES
VARIABLE CHARGES
Na3PO4
Groups IA,2A,3A,Zn, Cd, Ag
Transition metals, Sn, Pb
KHCO3
Rule:
Name of metal + name of anion
(or polyatomic ion)
Examples:
Ba(NO3)2
Fe2(SO4)3
MnO4Cr2O7
2-
CaF2
Cu2O
Al2O3
CuO
MgO
FeCl2
K2SO4
Fe3(PO4)2
NaCl
PbSO4
Comments: Ammonium ion,
NH4+, can act as the “metal”
Example:
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CHEMISTRY Rule:
Name of metal + (Roman
numeral to indicate ion’s
charge) + name of
anion/polyatomic ion.
Examples:
SnO2
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PRACTICE PROBLEMS
N2F4
BCl3
P2O5
SF4
ClF3
(COVALENT) Nonmetal + Nonmetal
WRITING FORMULAS
This system of writing formulas is based on whether the compound
had ionic or covalent bonds or is an acid.
The two common methods used are:
Least common multiple of charges
Crisscross method
RULES:
(1) Prefix + name of 1st element
a. No change in name
b.Mono is never used for the 1st
element.
(2) Prefix + name of 2nd element and change the
ending to –ide
a. Drop the “a” or “o” on the end of a
prefix if oxide is the 2nd element
EXAMPLES:
IONIC COMPOUNDS
Metal (or NH4+) + Nonmetal or Negative Polyatomic ion
Write the symbols for the ions side by side. Write the cation first.
Use parentheses for polyatomic ions: EX: (SO4)21. Write the symbols for the ions side by side. Write the cation
first. Use parentheses for polyatomic ions:
EX:
Iron (III) Oxide
CO2 –
N2O3 –
CHEMISTRY 2. Find the least common multiple of charges.
Iron (III) Oxide
PREFIXES:
1-
6-
2-
7-
3-
8-
4-
9-
5-
10 3. Find the number of ions needed to reach this multiple. (You can use
crisscross method as a shortcut.) The total positive charge must
match the total negative charge in the compound.
4. Get rid of the charges and then write the final formula
Iron (III) Oxide
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PRACTICE PROBLEMS
Barium sulfide
EXAMPLES
Dinitrogen trioxide (2 nitrogens and 3 oxygens), formula N2O3.
Phosphorus pentachloride (no prefix means 1 atom of phosphorus
with
5 chlorines), formula PCl5.
Calcium nitrate
PRACTICE PROBLEMS
Carbon dioxide
Ammonium sulfate
Phosphorous triiodide
We don’t bother writing subscripts for one (1) ion.
We ONLY use parentheses if BOTH:
It is a polyatomic ion; AND
We have more than one of them
Always reduce subscripts to the lowest ratio by dividing them by their
largest common factor.
PRACTICE PROBLEMS
Ammonium nitrate
Cobalt (II) sulfate
Nickel (II) cyanide
Barium oxide
Calcium hypochlorite
BINARY COVALENT COMPOUND
Nonmetal + Nonmetal
Prefix in front of each element tells you what subscript to use
Less electronegative element goes first and the second element
always ends in –ide.
Prefix mono is omitted for the first element in the name
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Sulfur dichloride
Dioxygen difluoride
Xenon trioxide
ACID NAMES & FORMULAS
H is the first element
Except for H2O and H2O2
Binary acid – contains only two different elements: Hydrogen and one
of the more electronegative elements
Oxyacids – acids that contain hydrogen, oxygen and a third element
(usually a nonmetal)
Acids are ionic formulas in which the positive ion is H+. Use as many
H+ ions as the charge on the negative ion.
EX: H+ and SO4 2- = H2SO4
Acids are named based on the name of the negative ion in the
compound and have the suffixes –ate, –ite, and –ide.
NOTE:
Two elements have variations.
H2SO4 – the sulfate ion changes to
H2SO3 – the sulfite ion changes to
H2 S
– the sulfide ion changes to
H3PO4 – the phosphate ion changes to
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SUFFIX OF NEGATIVE ION
1. Root name + ate
NAME OF ACID
Root name + ic acid
2. Root name + ite
Root name + ous acid
3. Root name + ide
THE MOLE
Mole was introduced by Wilhelm Ostwald in 1896
Mole comes from the Latin word “heap” or “pile”
1 mole = 6.022 x 1023 particles (Avogadro’s number)
Molecular Mass/Weight –Formula Mass/Weight
sum of the average atomic masses of all the atoms represented
in the formula of any molecule, formula unit or ion
# atoms x atomic mass
Hydro + Root name + ic acid
MOLAR MASS
Mass of one mole of a pure substance
Molar mass is numerically equal to formula mass
1 mole = Molar mass in grams
1 mole = 6.022 x 1023 particles (Avogadro’s Number)
Molar mass of a compound can be used as a conversion factor to
relate an amount in moles to a mass in grams for a given substance
PRACTICE PROBLEMS
H2C2O4
HF
Perchloric acid
Carbonic acid
Sulfurous acid
COMMON NAMES
PAINS IN THE GLUTEUS MAXIMUS: these lovely creatures have been
around longer than the naming system and no one wanted to adapt!!
Water
Ammonia
Hydrazine
Phosphine
Nitric oxide
Nitrous oxide
(laughing gas)
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PRACTICE PROBLEMS
Calculate the molar mass of
a) Ca3(PO4)2
b) caffeine, C8H10N4O2
c) BaCl2•2H2O
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MASS-MOLE-MOLECULES CONVERSIONS
Moles to Grams
4.27 mol TiO2 to grams
Molecules or Formula Units to Moles
4.36 x 1025 molecules C6H5CH3 to moles
Grams to Molecules
Grams to Moles
50.0 g NH3 to moles
72.5 grams CHCl3 to molecules
Moles to Molecules or Formula Units
Molecules to Grams
0.30 mol Sr(NO3)2 to formula units
8.39 x 1023 molecules HF to grams
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DETERMINATION OF THE FORMULA OF A COMPOUND
PERCENT COMPOSITION
Law of constant composition states that any sample of a pure
compound always consists of the same elements combined in the
same proportions by mass.
mass of element
If mass in grams is given:
%
x 100
total mass
If only the formula of the compound is given:
mass of element
%
x 100
molar mass of compound
A good check is to see if the results add up to 100% (rounding will
cause a slight discrepancy)
Your final answers should have 2 decimal places
PRACTICE PROBLEM:
Express the composition of each of the following compounds in terms of
the mass of each element in 1.00 mole of compound and the weight
percent of each element:
a) Ba3(PO4)2
DETERMINATION OF EMPIRICAL FORMULAS
Consist of the symbols for the elements combined in a compound,
with subscripts showing the smallest whole-number ratio of the
different atoms in the compound
If % composition is given, assume 100 g of substance
% = mass in grams of element
Calculate # mole of each element (mass/atomic mass)
Divide each # of mole by the smallest value to get the simplest ratio
of subscripts.
Multiply by an integer if needed to convert to whole number
.5 x 2, .33 x 3, .25 x 4
Determine the empirical formula of a compound containing 52.11% C,
13.14% H and 34.75% O
A 175 g sample of compound contains 56.15 g C, 9.43 g H, 74.81 g O,
13.11 g N, and 21.49 g Na. Find the empirical formula.
b) MgSO47H2O
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Determine the empirical formula of a compound containing 40.9% carbon,
4.58% H and 54.5% O.
DETERMINATION OF MOLECULAR FORMULA
The actual formula of a molecular compound; gives actual # of atoms
present in the compound
Find empirical formula (EF) if it is not given
Calculate EF Mass (same as molar mass calculation)
Determine the multiple (n) – actual ratio of elements in compound
5. MF = (EF)n
n
Molecular Formula mass
Empirical Formula mass
What is the
molecular formula
of a compound that has an empirical formula of CH2O and a molar mass
of 120.12 g/mol?
The compound borazine consists od 40.29% boron, 7.51% hydrogen,
and 52.20% nitrogen, and its molar mass is 80.50 g/mol. Determine
the molecular formula of borazine.
HYDRATES
Formula units with water associated with them.
The water molecules are incorporated into the solid structure.
EX: CuSO45H2O, copper (II) sulfate pentahydrate.
Strong heating can generally drive off the water in these salts.
Once the water has been removed the salts are said to be anhydrous
(without water).
PRACTICE PROBLEM
Cerium (III) iodide occurs as a hydrate with the composition 76.3%
CeI3 and 23.7 % water. Determine the formula for the hydrate.
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