Oxidation Numbers & Formulas• Matter & its

states• Laws of

thermodynamics• Measuring &

Calculating• Atomic Structure• Elements – the

Periodic Table• Chemical Bonds

Oxidation Numbers and FormulasChemical

Composition and Reactions◦Valence bonding◦Bookkeeping

system Electrons involved

in bonding Oxidation numbers Assign each

electron to a element in compound

Oxidation Numbers and FormulasOxidation 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 RuleElements in their natural state (pure

elements) = 0Also applies to Mr. H. BrClFONDiatomic Elements share electrons equally

Oxidation Numbers and Formulas2. The Ion Rule

◦ The oxidation # of a monatomic ion is equal to the charge of the ion

Br- = -1 Mg2+ = 2

Oxidation Numbers and Formulas3. The zero sum

rule The sum of the

#s in a compound must be zero

◦ 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 Assigned – ox. number

◦ Lower EN element “loses” electrons

Assigned + ox. Number◦ Element with highest EN usually

determines ox. #’s of other elements

Oxidation Numbers and Formulas

4. A. Alkali metals always have a +1 oxidation numberB. Alkaline earth metals always have a +2 ox. #C. Hydrogen usually has a +1 when bonded with nonmetals, -1 when bonded with metalsD. Oxygen always has a -2 except when bonded with fluorine (+2 – Fl has higher EN so it takes the electrons)Peroxide ion O2

2- Oxygen has a -1E. Halogens = -1 when bonded to metalsBonded 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!

Oxidation Numbers and Formulas

Rule Summary1. Free atoms = 02. Ion charge = ox. #3. Compound sum = 04. A. Group 1 = +1

B. Group 2 = +2C. H = +1 or -1D. O = -2 or -1E. Group 17 (halogens) = -1

5. 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 Tri-

Chloride 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 crysta

lline 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

Hydrogen H + Sulfate SO42- = Sulfuric acid H2SO4

◦ Anion ends in –ite, ending changes to –ous + acid

Hydrogen H + Sulfite SO32- = Sulfurous acid H2SO

3

◦ 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

Hydrogen H + Sulfate SO42- = Sulfuric acid H2SO4

◦ Anion ends in –ite, ending changes to –ous + acid

Hydrogen H + Sulfite SO32- = Sulfurous acid H2SO

3

◦ 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 + O2 H2O

What Equations Do H2 + O2 = H2O Must be same amount of

atoms on left as on right◦ 1st law of thermodynamics

So must balance it H2 + O2 = H2O H’s are balanced, O’s are

not Double H2O’s

◦ H2 +O2= 2H2O Now H’s unbalanced

◦ Double H on left◦ 2 H2 +O2= 2H2O

Now balanced Going back and forth

normal

2 H2 +O2= 2H2OBalanced Chemical

EquationProcess 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

A + BC AC + B ◦ Double replacement/displacement

You swap AB + CD AC + BD◦ Classes of reactions

Single Replacement Reactions More active vs. less active metals

◦ Usually form precipitates Reactions in acids

◦ Replace hydrogen which bubbles out Reactions in water

◦ Alkali metals – hydrogen bubbles out Halogen to halogen in solution

◦ 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