1 / 38

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

carnig
Télécharger la présentation

Oxidation Numbers & Formulas

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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

More Related