Oxidation Numbers & Formulas
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Matter & its states
Laws of
thermodynamics
Measuring &
Calculating
Atomic Structure
Elements – the
Periodic Table
Chemical Bonds
Oxidation Numbers and Formulas

Chemical
Composition and
Reactions
◦ Valence bonding
◦ Bookkeeping system
 Electrons involved in
bonding
 Oxidation numbers
 Assign each electron to
a element in compound
Oxidation Numbers and Formulas

Oxidation Number
◦ # of electrons that an
atom in a compound
must gain or lose to
return to its neutral
state.
 Neg. number – element
has gained that many
electrons
 -2=how many?
 Pos. number – element has
given up that many
electrons
 +2 = how many?
Oxidation Numbers and Formulas

#s Originally assigned based up
experimentation
◦ Analysis to determine chemical composition

Now Use Rules
◦ Predict how elements typically combine
◦ Of course there are exceptions
1.
The Free Element Rule
Elements in their natural state (pure elements) = 0
Also applies to Mr. H. BrClFON
Diatomic Elements share electrons equally
Oxidation Numbers and Formulas
The Ion Rule
2.
◦
The oxidation # of a
monatomic ion is equal
to the charge of the
ion


Br- = -1
Mg2+ = 2
Oxidation Numbers and Formulas
The zero sum rule
The sum of the #s in
a compound must be
zero
3.

◦
Compounds are not
electrically charged
Oxidation Numbers and Formulas
Ionic compounds

◦
◦
◦
NaCl (+1/-1)
MgCl2 (+2/-1(2))
Formula unit perfectly
balances the charges
Oxidation Numbers and Formulas
Covalent Compounds

◦
Shared electrons closer to higher EN
element in compound

◦
Lower EN element “loses” electrons

◦
Assigned – ox. number
Assigned + ox. Number
Element with highest EN usually
determines ox. #’s of other elements
Oxidation Numbers and Formulas
A. Alkali metals always have a +1 oxidation number
B. Alkaline earth metals always have a +2 ox. #
C. Hydrogen usually has a +1 when bonded with
nonmetals, -1 when bonded with metals
D. Oxygen always has a -2 except when bonded with
fluorine (+2 – Fl has higher EN so it takes the
electrons)
Peroxide ion O22- Oxygen has a -1
E. Halogens = -1 when bonded to metals
Bonded to nonmetals, element with higher EN
assigned negative number. Fl always -1 since it
has highest electronegativity
5.
Sum of oxidation #s in a polyatomic ion = charge of
ion
If rules contradict each other, closer to 0 rule rules!
4.
Oxidation Numbers and Formulas

1.
2.
3.
4.
5.
Rule Summary
Free atoms = 0
Ion charge = ox. #
Compound sum = 0
A. Group 1 = +1
B. Group 2 = +2
C. H = +1 or -1
D. O = -2 or -1
E. Group 17 (halogens) = -1
Sum of Ons in a polytamic ion = charge
Multiple Oxidation States

◦
Some atoms have
multiple
Depends on other
elements bonding





Especially trans metals
Outer energy levels close
proximity
d & f sublevels
Depends on # of electrons
participating in bonding -=
FeCl2 FeCl3
Memorize ‘em or look ‘em up
Some nonmetals multiple
too



N = 5 to -3
ON driven by higher EN
element!
Polyatomic Ions
Covalently bonded
atoms that carry a
charge

◦
◦
Own rule
ON of atoms in a poly
ion add up to its charge
OH- ON’s: O=-2, H=1
◦


Sum = -1, its charge
Poly ions survive most
chemical reactions intact,
so treat as separate ON,
just like an element
Nomenclature
Times past – given
common name

◦
Associates with
compound – place
mined or some
characteristic

◦
Milk of magnesia, etc.
Tell nothing about
composition or formula

Table 8-2
Nomenclature

◦
More and more
compounds discovered,
realized must have
reliable naming system
IUPAC developed
standardized set of rules
call nomenclature



Which elements present,
type of compound,
intermolecular attractions,
general properties
Soda ash – sodium
carbonate – Na2CO3
Epsomite – Magnesium
sulfate – Mg(SO4)2
Binary Covalent Compounds
Binary Covalent
Compounds

◦
◦
◦
◦
Two elements, bonded
covalently
Acids – begin with
hydrogen (usually)
HCl – hydrochloric acid
– in your gut and your
pool
H2SO4 – sulfuric acid –
in your car battery
Binary Covalent Compounds
Greek Prefix System

◦
How many of each in a
covalent compound
Table 8-3
Mono used for second
element (unless needed
for clarity) – extra
vowels eliminated
◦
◦

◦
◦
Carbon monoxide non
mono-oxide
Least EN element first
Ending of last element
changed to -ide
Binary Covalent Compounds
Flow Chart 8-4
HCl


◦

◦

◦

◦
Acid? Acid rules (8-12)
PCl3
Phosphorus TriChloride
CO2
Carbon Dioxide
H2O
Dihydrogen Monoxide
Binary Ionic Compounds
Not named using
Greek prefix system

◦
2 element compounds
Metal – Nonmetal
Named after 2 ions
involved
Cation – Element name




E.g. Sodium
Anion –ide ending



Chlorine becomes Chloride
Sodium Chloride
Binary Ionic Compounds
Polyatomic Ionic Compounds
Ions with multiple elements (2 or more)

◦
A compound with a charge
Of common ions, only positive (cations) are
ammonium NH4 and the mercurous ion Hg22+
All the rest anions
Ions containing oxygen and one other called
oxyanions
Number of oxygen atoms drives the name
Often 2 or more forms perchlorate, chlorate, chlorite
and hypochlorite – all chlorine and oxygen





◦
Bromide family same way – usually halogens
If only two ions, fewer oxygens is _ite, more _ate

◦
Sulfite, sulfate
Naming Polyatomic Ionic Compounds
Simple – just name the
cation and anion, just
as with binary ionics

◦
Table 8-8
Ion generally comes last
since only 2 common
cations


◦
But if first – notice the _ide
ending just as with binaries
Example problems 8-7,
8-8
Ionic Compounds and Multiple
Oxidation States
Metal in ionic compound have
more than 1 oxidation state?

◦
◦

Roman numeral after name to
show ON
Stock or Roman numeral system
Flow chart and ex. Problem 8-9,
8-10
Hydrates
Compounds that hold a
characteristic amount of water in
their crystalline structure

◦
Water of hydration
Combine in specific ratios due to crystalline
structure
Formulas indicate water with dot #H2O





E.g. (Na2CO3 . 7H2O) – Sodium carbonate
heptahydrate
No water present? Anhydrous
See table 8-10
Binary Acids


◦
Covalent compounds
usually beginning with
hydrogen
H + 1NM= binary acid
When liquid, different
naming scheme
HCl – when gas—
hydrogen chloride
◦

◦
Dissolved in water—
hydrochloric acid
Naming – hydro + NM
root name + ic acid

HBr becomes Hydrobromic
acid
Acid Burns
Ternary Acids
3 elements – H, O and another NM
O and NM often a polyatomic ion
Names derived from anions in acid



◦
Anion ends in –ate, ending changes to –ic +
acid

◦
Anion ends in –ite, ending changes to –ous +
acid

◦
◦
Hydrogen H + Sulfate SO42- = Sulfuric acid H2SO4
Hydrogen H + Sulfite SO32- = Sulfurous acid
H2SO3
Table 8-12
Ex. Problem 8-11
Ternary Acids
3 elements – H, O and another NM
O and NM often a polyatomic ion
Names derived from anions in acid



◦
Anion ends in –ate, ending changes to –ic +
acid

◦
Anion ends in –ite, ending changes to –ous +
acid

◦
◦
Hydrogen H + Sulfate SO42- = Sulfuric acid H2SO4
Hydrogen H + Sulfite SO32- = Sulfurous acid
H2SO3
Table 8-12
Ex. Problem 8-11
Writing Equations
Visible signs of unseen chemical
reactions that hint at molecular
change

◦
◦
◦
Bubbles in pancakes/biscuits
One chemical combines with another to
create a new substance
Scientists call these changes Chemical
Reactions
What reacted? What was produced?
How much of each?

◦
Answers in a balanced chemical equation
What Equations Do
Describe chemical reactions

◦
◦
◦
◦
ID all substances in a reaction
Left side=reactants
Right side=products
Word equation – all substances but
not quantities

◦
◦
Hydrogen + Oxygen
Water
Formulas show quantity and
composition
H2 + O 2
H 2O
What Equations Do
H 2 + O2 = H 2 O
Must be same amount of
atoms on left as on right


◦




◦
So must balance it
H 2 + O2 = H 2 O
H’s are balanced, O’s are
not
Double H2O’s
H2 +O2= 2H2O
Now H’s unbalanced

◦
◦


1st law of thermodynamics
Double H on left
2 H2 +O2= 2H2O
Now balanced
Going back and forth
normal
2 H2 +O2= 2H2O
Balanced Chemical Equation
Process called:
Balancing by inspection
What Equations Do






Look at one on pp. 196-7
Calcium hydrogen carbonate + calcium
hydroxide yields water + calcium carbonate
Ca(HCO3)2 + Ca(OH)2
H2O + CaCO3
2Ca, 4H, 2C, 8O
2H, 1Ca, 1C, 4O
Everything on right exactly ½ of left
Ca(HCO3)2 + Ca(OH)2 2H2O + 2CaCO3
Balancing by Inspection
BOTH SIDES MUST BALANCE!

◦
◦
◦
◦
◦
◦
Equal numbers of each atom on both
sides
Nitrogen monoxide +oxygen
nitrogen
dioxide
NO + O2
NO2
1N, 3O’s
1N, 2O’s
NO FRACTIONS
Must be in lowest terms
Balancing by Inspection

◦
◦
◦
◦
◦
Reversible Reactions
Can happen both ways
Gas (g) or
Liquids (l)
Solid (s) or
Dissolved in water – aqueous (aq)
All acids are aqueous
H2SO4 – hydrogen sulfate
H2SO4(aq) Sulfuric acid
Solid falls out of solution – precipitate




◦


Precipitation sometimes noted with
See ex. On p. 198
Table 8-13 – more symbols
Limitations of Equations
Cannot predict if a reaction
will occur
Do not tell if equation will go
to completion


◦
◦
Some take several steps
Chemical reactions
Reactions/Relationships
Synthesis reaction – A +B

◦
AB
You “go out” with a single

Examples in book, pp. 203-4
Decomposition reaction

◦
You breakup – AB

A+B
Examples in book, p. 205
Replacement/Displacement reactions

◦
Single replacement; You replace somebody else

◦
AC + B
Double replacement/displacement

◦
A + BC
You swap AB + CD
Classes of reactions
AC + BD
Single Replacement Reactions
More active vs. less active metals

◦
Reactions in acids

◦
Replace hydrogen which bubbles out
Reactions in water

◦
Alkali metals – hydrogen bubbles out
Halogen to halogen in solution

◦

Usually form precipitates
More vs. less reactive
Activity series allows prediction
Replacement Reactions
Double Replacement Reactions
Aqueous mixtures of 2 ionic compounds

◦
◦
◦
Precipitate forms – evidence of reaction
Solution breaks ions apart, allows reaction
Ionic equation – only for reactions in solution
All particles present before and after solution
Insoluble ions represented by (s)
Include particle not participating






Spectator ions
Stricken from equation
Net Ionic Equation

Only ions reacting

Example – p. 206
Double Replacement Reactions
◦
Neutralization reactions
HCl + KOH

HOH + KCl

H+(aq) + Cl- (aq) + K+(aq) + OH-(aq)
HOH(l) + K+(aq) + Cl-(aq)


Cl- and K+ are spectator ions
All neutralization reactions have
same net ionic equation



Water created
Easy to separate salt since water can
be boiled away
Double replacement reations usually
reduce # of ions in solution