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Oxidation Numbers & Formulas

Oxidation Numbers & Formulas. 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

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Oxidation Numbers & Formulas

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  1. Oxidation Numbers & Formulas • Matter & its states • Laws of thermodynamics • Measuring & Calculating • Atomic Structure • Elements – the Periodic Table • Chemical Bonds

  2. 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

  3. 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?

  4. 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 • The Free Element Rule Elements in their natural state (pure elements) = 0 Also applies to Mr. H. BrClFON Diatomic Elements share electrons equally

  5. Oxidation Numbers and Formulas • The Ion Rule • The oxidation # of a monatomic ion is equal to the charge of the ion • Br-= -1 • Mg2+ = 2

  6. Oxidation Numbers and Formulas • The zero sum rule • The sum of the #s in a compound must be zero • Compounds are not electrically charged

  7. Oxidation Numbers and Formulas • Ionic compounds • NaCl (+1/-1) • MgCl2 (+2/-1(2)) • Formula unit perfectly balances the charges

  8. 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

  9. 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 • Sum of oxidation #s in a polyatomic ion = charge of ion If rules contradict each other, closer to 0 rule rules!

  10. Oxidation Numbers and Formulas • 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

  11. 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!

  12. 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

  13. 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

  14. 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

  15. 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

  16. 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

  17. Binary Covalent Compounds • Flow Chart 8-4 • HCl • Acid? Acid rules (8-12) • PCl3 • Phosphorus Tri-Chloride • CO2 • Carbon Dioxide • H2O • Dihydrogen Monoxide

  18. 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

  19. Binary Ionic Compounds

  20. 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

  21. 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

  22. 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

  23. 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

  24. 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

  25. 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 H2SO3 • Table 8-12 • Ex. Problem 8-11

  26. 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 H2SO3 • Table 8-12 • Ex. Problem 8-11

  27. 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

  28. 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

  29. 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= 2H2O • Balanced Chemical Equation • Process called: • Balancing by inspection

  30. 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

  31. 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

  32. 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

  33. 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

  34. 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

  35. 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

  36. Replacement Reactions

  37. 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

  38. 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

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