oxidation numbers & formulas matter & its states laws of thermodynamics measuring &...
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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 O22- Oxygen has a -1
E. 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 cryst
alline 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 H2
SO3
◦ 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 H2
SO3
◦ 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