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CHEMISTRY: about 250 years old as a science Ancient roots Ancient roots II. medical roots common trades: mining and metalworking physiological effects of certain substances finding or making medicines glassmaking food processing (wine making, vinegar production) alchemy ال soap making origin of word 'chemistry' (.... chimie, Chemie, химия) Arabic 'al' ~ 'the' Middle Ages: transformation of ordinary metals into gold Chema: ancient province in Egypt famous for its Sir Isaac Newton (1643-1727) was also an (unsuccessful) alchemist soap production or kēme (chem) = „earth” in Egyptian Greek science Greek science Democritus (460370 BC) Aristotle (384322 BC) philosopher philosopher and experimental scientist atomic theory '' = indivisble discrete theory of matter continuous matter fundamental elements: earth, air, water, fire fundamental properties: cold, hot, dry, wet Definition of a chemical element Robert Boyle (1627-1691) Examples of elements and non-elements Iron: element water: not an element, compound gold: element air: not an element, mixture of elements and compounds An element is a fundamental substance that cannot be chemically changed or broken down into anything simpler. STILL CORRECT AS MODERN DEFINITION brass or bronze: mixture of elements rows = periods columns = groups noble gases alkali earth metals Chemistry as a quantitative science halogens end of 18th century: Lavoisier and Lomonosov Antoine Laurent Lavoisier (1743-1794) Mikhail Vasilyevich Lomonosov (1711-1765) alkali metals Chemistry as a quantitative science Physical properties end of 18th century: Lavoisier and Lomonosov extensive intensive accurate measurements (mass, volume, ...) related to size e.g. mass unrelated to size e. g. temperature discovery of gases Quantitative measurements fundamental not a combination of other properties e.g. length V = 330 ml derived a combination of other properties e.g. area 1960: Systéme Internationale d'Unites (SI) physical property number unit (volume) (milliliter) (330) earlier systems: e.g. CGS, MKS Fundamental properties and units in SI property Prefixes in SI symbol name of unit symbol of unit m kilogramm kg mass length meter l m temperature T kelvin K amount of substance n mole mol time t second s electric current I ampere A luminous intensity IV candela cd prefix symbol factor prefix symbol factor giga G 109 deci d 101 mega M 106 centi c 102 kilo k 103 milli m 103 hecto h 102 micro 106 deka da 101 nano n 109 e.g. 1 nm = 109 meter MOST IMPORTANT DERIVED PROPERTIES Volume: m3 CONVERSION OF UNITS Volume: dm3 = L (or l) cm3 = mL (or ml) l (or L) Density: kg/m3 1 kmol = 103 mol m3 1 dm3 = 0.001 m3 = 1000 cm3 = 106 mm3 Density: g/cm3 kg/m3 Pressure: Pa (pascal) = kgm-1s-2 Pressure: atm, bar, torr Energy: J (joule) = kgm2s-2 1 g/cm3 = 0.001 kg/ 106 m3 = 1000 kg/m3 Pa (pascal) = kgm-1s-2 1 hPa = 100 Pa 1 MPa = 106 Pa Energy: cal J (joule) = kgm2s-2 1 mJ = 0.001 J 1 nJ = 109 J Significant figures (s.f.): the total number of digits in a measurement Accuracy and Precision: Correct use: Accuracy: shows how close to the true value a given measurement is Precision: shows how well a number of independent measurements agree with one another all digits but the last are certain, the last is a best guess, usually having an error ± 1 4.803 cm 4 0.00661 g 3 55.220 K 5 34200 m 3-5 SOME NUMBERS ARE EXACT!! Significant figures (s.f.): Significant figures and calculations all digits other than 0 significant 1. Zeros in the middle of a number are always significant 2. Zeros at the beginning of a number are never significant 3. Zeros at the end of the number and after the decimal point are always significant 4. Zeros at the end of a number and before the decimal point may or may not be significant 1. Multiplication or division: the answer cannot have more significant numbers than either of the original numbers 11.78945 g / 11.9 cm3 = 0.991 g/cm3 2. Addition or subtraction: the answer cannot have more digits to the right of the decimal point than either of the original numbers 3.18 g + 0.01315 g = 3.19 g Rounding numbers 5.664525 1. If the first digit you remove is less than 5, round by dropping it and all following digits. (3 s.f.) 5.66 2. If the first digit you remove is 6 or greater, round by adding 1 to the digit on the left (2 s.f.) 5.7 3. If the first digit you remove is 5 and there are more nonzero digits following, round up. (4 s.f.) 5.665 4. If the digit you remove is 5 with nothing following, round down. (6 s.f.) 5.66452 LAW OF MASS CONSERVATION Mass is neither created nor destroyed in chemical reactions. TODAY: Not exactly true Einstein's mass-energy equivalence equation DE = Dmc2 ~100 kJ/mol ~10-12 kg/mol = 1 ng/mol LAW OF DEFINITE PROPORTIONS 1799 Joseph Louis Proust (1754-1826) LAW OF MULTIPLE PROPORTIONS John Dalton (1766-1844) 1803 Different samples of a pure chemical substance always contain the same proportion of elements by mass. 1.000 g Cu 400 °C 1.252 g CuO 400 °C HNO3 Cu(NO3)2 Elements can combine in different ways to form different substances, whose mass ratios are small whole-number multiples of each other. Example: nitrogen oxides Atomic theory: compound %N %O mO/mN ratio N2 O 63.7 36.3 0.571 1 NO 46.7 53.3 1.142 2 N2 O 3 36.9 63.1 1.713 3 NO2 30.4 69.4 2.284 4 N2 O 5 25.9 74.1 2.855 5 John Dalton, 1808 • Elements are made of tiny particles called atoms. • Each element is characterized by the mass of its atoms. Atoms of the same element have the same mass, but atoms of different elements have different masses. • Chemical combination of elements to make different substances occurs when atoms join together in small whole-number ratios. • Chemical reactions only rearrange the way that atoms are combined; the atoms themselves are unchanged. LAW OF COMBINING GAS VOLUMES Most elements: atoms and bonds between atoms in a well defined spatial arrangement 1808 Joseph Louis Gay-Lussac (1778-1850) Some elements: small independent units containing a few identical atoms MOLECULES H2, N2, O2, F2, P4, S8, Cl2, Br2, I2 AVOGADRO'S LAW Gases at constant temperature and pressure react or are produced in a chemical reaction in volume ratios of the small whole numbers. HOW SMALL IS AN ATOM? WHAT IS THE MASS OF AN ATOM? Amedeo Avogadro (1776-1856) 1811 Identical volumes of different gases contain the same number of molecules at the same temperature and pressure. first: relative atomic masses atomic mass unit (amu): exactly one twelfth of the mass of a carbon-12 atom (1.660539 1027 kg) later: numbers of atoms or molecules are conveniently measured by the mole 1 mol: the number of atoms in exactly 12 g carbon-12 Avogadro's number NA = 6.022 1023 mol-1 ATOMIC NUMBER (Z) = STRUCTURE OF ATOMS Number of protons in atom's nucleus = Number of electrons around atom's nucleus electron cloud: electrons MASS NUMBER (A) = protons nucleus Number of protons (Z) + Number of neutrons (N) neutrons Hydrogen: mass (amu) electric charge electron 0.00055 1e proton 1.00728 +1e neutron 1.00866 0 mass number e = 1.602 1019 C C 1 proton + 1 neutron (deuterium) 1 proton + 2 neutrons (tritium) Isotopes = atoms with identical atomic numbers but different mass numbers Many elements: mixtures of isotopes. symbol 12 6 1 proton + O neutron (protium) carbon - 12 Chlorine: atomic number 75.77% 35Cl 34.969 amu 37Cl 36.966 amu 24.23% 31 15 P phosphorus - 31 Average atomic mass of chlorine: 235 92 268 109 U Mt uranium - 235 meitnerium - 268 0.7577 34.969 amu + 0.2423 36.966 amu = 35.46 amu The isotopic composition of an element is constant in nature. Elements Compound: a pure substance that is formed when atoms of two or more different elements combine and create a new material with properties completely unlike those of its constituent elements. Chemical compounds element Pure Substances Matter @ symbol compound @ chemical formula Heterogeneous Mixtures NaCl, H2O, C12H22O11, ThO2, BaTiO3, Ca(ClO4)2 Chemical equation: Homogeneous 2 H2 + O2 2 H2O reactants products Mixture: a simple blend of two or more substances added together in some random proportion without chemically changing the individual substances themselves. Chemical changes: atoms interact primarily through the electron structure Covalent bond: atoms share electrons, bonding electrons belong to each atom forming the bond Heterogeneous: the mixing is not uniform, the mixture has regions of different composition Molecule: the unit of matter that results when two or more atoms are joined by covalent bonds example: water + oil O2, N2, F2, I2, P4, S8 Homogeneous: the mixing is uniform example: dust-free air H2O, NH3, HF, C3H8 C2H6O, C2H5NO2 Ionic bond: transfer of one or more electrons from one atom to another Metallic bond: sharing of electrons between a large number of metal atoms Na + Cl Na+ + Cl cation Metallic solid: metal atoms connected by metallic bonding and packed together in a regular way anion Ionic solid: anions and cations packed together in a regular way Na, Ti, Fe, Cu, In, Pb NaCl, CsCl, MgCl2, CaO, Ce(SO4)2 Elements: Symbols (Ca for Calcium, Sn for Tin, .....) Compounds: Formulas NaCl formula for sodium chloride or common salt chemical name systematic name common name trivial name Formulas: different types for different purposes Formulas empirical formula: shows the ratios of atoms in a compound NaCl CH2O molecular formula: shows the actual number of atoms in a molecule ----CH2O (CH2O)6 = C6H12O6 structural formula: shows how the atoms are connected CHO within a molecule H OH H OH HO H HO H ----HO C O H OH OH HO H H OH H OH H H CH2OH Example: NAMING COMPOUNDS unknown compound 92.26 % C and 7.74% H 100 g compound: 92.26 g C 92.26 / 12.011 = 7.681 mol C only inorganic compounds, organic compounds will be covered as part of the Organic Chemistry course 7.74 g H 7.74 / 1.008 = 7.68 mol H empirical formula: C7.681H7.681 CH if M = 26 g/mol C 2H 2 (acetylene) if M = 78 g/mol C 6H 6 (e.g. benzene) NOMENCLATURE = system of names, naming rules present chemical nomenclature: discussed and approved by IUPAC percent composition: (experimental) elemental analysis International Union of Pure and Applied Chemistry molecular mass: (experimental) mass spectrometry BINARY IONIC COMPOUNDS BINARY IONIC COMPOUNDS CrCl3: chromium(III) chloride name = positive ion + negative ion Cr can form both Cr2+ and Cr3+ ions! positive ion = name of the element PbS: negative ion = name of the element + suffix -ide lead(II) sulfide FeCl2: iron(II) chloride ferrous chloride KCl: potassium chloride FeCl3: iron(III) chloride ferric chloride LiF: lithium fluoride SnCl2: tin(II) chloride stannous chloride SnCl4: tin(IV) chloride stannic chloride BaCl2: barium chloride Cu2O: copper(I) oxide cuprous oxide Ba always forms Ba2+ ions AlBr3: aluminum bromide CuO: copper(II) oxide cupric oxide Al always forms Al3+ ions BINARY MOLECULAR COMPOUNDS name = more positive part + more negative part (suffix –ide) use of Greek numbers in names: BINARY MOLECULAR COMPOUNDS N2O: dinitrogen monoxide HF: hydrogen fluoride CO: carbon monoxide CO2: carbon dioxide 1 --- (mono) 5 penta 2 di 6 hexa SeBr4: selenium tetrabromide 3 tri 7 hepta PCl5: phosphorus pentachloride 4 tetra 8 octa AsCl3: arsenic trichloride SF6: sulfur hexafluoride IF7: iodine heptafluoride NO: nitrogen monoxide N2O3: dinitrogen trioxide NO2: nitrogen dioxide N2O4: dinitrogen tetroxide N2O5: dinitrogen pentoxide SPECIAL CASES POLYATOMIC IONS names must be learnt e.g. CO32- carbonate ion SO42- sulfate ion OsO4: osmium tetroxide or osmium(VIII) oxide NO3- PO43- phosphate ion CrO3: chromium trioxide or chromium(VI) oxide NH4+ ammonium ion Some metal compounds (often oxides) are molecular nitrate ion MnO2: manganese dioxide or manganese(IV) oxide (NH4)2SO4 ammonium sulfate Very rare exceptions: Fe(NO3)2 iron(II) nitrate or ferrous nitrate V2O5: vanadium pentoxide or vanadium(V) oxide Hg(NO3)2 mercury(II) nitrate or mercuric nitrate P2O5: phosphorus pentoxide KCN potassium cyanide (CN- cyanide ion) Balancing chemical equations Chemical equations conservation of mass and conservation of charge molecular equation: 3CaCl2 + 2Na3PO4 Ca3(PO4)2 + 6NaCl AgNO3 + NaCl AgCl + NaNO3 ionic equation (aqueous phase = in water): Ca: 3 = Cl: 3×2 = 6 = 2×3 = 6 = 3 6 Ag+ + NO3 + Na+ + Cl AgCl + Na+ + NO3 Na: P: 2 = 2 6 Ag+ + Cl AgCl O: 2×4 = 8 = 2×4 = 8 charge: 0 = 0 Balancing chemical equations Balancing chemical equations Step 1: Write reactants and products C3H8 + O2 S: 2S2O32 + I2 2×2 = 4 = O: I: 2×3 = 6 2 = = S4O62 + 2I 4 C3H8 + O2 6 2×1 = 2 2×-2 = -4 = -2+2×-1 = -4 CO2 + H2O CO2 + H2O Step 3: Find coefficients to balance this atom 1C3H8 + O2 charge: Step 2: Find one atom that occurs only in one substance on both sides 3CO2 + H2O Step 4: Find another unbalanced atom which occurs in only one substance 1C3H8 + O2 3CO2 + H2O Stoichiometry Step 5: Find coefficient to balance this atom 1C3H8 + O2 3CO2 + 4H2O (Greek) στοιχεῖον = element and μέτρον = measure Step 6: Repeat steps 4-5 until you have balanced all atoms 1C3H8 + O2 3CO2 + 4H2O 1C3H8 + 5O2 3CO2 + 4H2O quantitative relationship between reacting or produced substances (mass, amount of substance) C2H4 (g) + HCl (g) C2H5Cl (g) Step 7: Make sure the coefficients are reduced to the smallest whole number values 1C3H8 + 5O2 1 mol hydrogen chloride to produce 3CO2 + 4H2O 1 mol ethyl chloride Step 8: Check the balanced equation 1C3H8 + 5O2 3C, 8H, 10O gas 1 mol ethylene reacts with 28.02 g ethylene reacts with 3CO2 + 4H2O 36.46 g hydrogen chloride to produce 3C, 10O, 8H 64.48 g ethyl chloride Limiting amounts of reactants C2H4 (g) + HCl (g) C2H5Cl (g) How much hydrogen chloride reacts with 15.00 g ethylene? 15.00 g ethylene = 15.00g/28.02gmol-1 = 0.5353 mol C2H4 (g) + HCl (g) C2H5Cl (g) What happens if we have 5 mol C2H4 and 3 mol HCl at the beginning? 5 mol C2H4 would react with 5 mol HCl HCl not enough 0.5353 mol ethylene reacts with 3 mol HCl would react with 3 mol C2H4 C2H4 too much 0.5353 mol hydrogen chloride After reaction: 3 mol C2H5Cl and 2 mol C2H4 unchanged 0.5353 mol hydrogen chloride = HCl: limiting reagent, completely consumed 0.5353 mol × 36.46g/mol = 19.51 g C2H4: excess reagent, some remains unreacted Yields of chemical reactions C4H8 (g) + CH4O (l) C5H12O (l) Stoichiometry: calculated for an IDEAL case when the limiting reagent is completely converted to the desired product Reality: 1. products other than the desired one may form (side reactions) 2. some of the limiting reagent may remain unchanged YIELD: the amount of desired product actually formed Actual yield of desired product PERCENT YIELD = Theoretical yield of desired product liquid From 26.3 g of isobutylene and excess methanol 32.8 g of methyl tert-butyl ether is prepared. What is the percent yield? 26.3 g C4H8 (56.12 g/mol) 0.469 mol Theoretical yield of C5H12O: 0.469 mol × 88.17 g/mol = 41.35 g PERCENT YIELD = 32.8 g / 41.35 g = 79.3 % × 100% A chemical equation HCl (g) in gas phase does show - identity of reactants and products - stoichiometry CH4O (l) in liquid phase - direction of spontaneous reaction 2H2 + O2 2H2O SiO2 (s) in solid phase may show NaCl (aq) in aqueous phase = dissolved in water or 2H2 + O2 = 2H2O - phase of reactants and products - energy change of the reaction 2H2 (g) + O2 (g) 2H2O (g) DE = 484 kJ/mol does not usually show CaCO3 (s) + 2CH3COOH (l) - time needed for completion Ca(CH3COO)2 (aq) + H2O (l) + CO2 (g) - conditions (e.g. temperature) EXAMPLE: SOLUTION Coca-Cola = SOLUTE(S) + SOLVENT carbon dioxide water sugar 5.0 g glucose dissolved in water to make 100 mL solution. 5.0 g C6H12O6 / 180.18 g/mol = 0.028 mol molarity: 0.028 mol / 0.100 L = 0.28 M (many) other substances AgNO3 + NaCl AgCl + NaNO3 concentration various units Equivalent quantities: reacting exactly with each other molarity (M): most important in chemistry 1.0 L of 1.0 M AgNO3 and 1.0 L of 1.0 M NaCl 1.0 L of 1.0 M AgNO3 and 2.0 L of 0.5 M NaCl Moles of solute MOLARITY = Volume of solution unit: molL-1 mol/dm3 M DILUTION: 4.0 L of 0.25 M AgNO3 and 1.0 L of 1.0 M NaCl 0.40 L of 0.10 M AgNO3 and 0.20 L of 0.20 M NaCl GROUPS OF CHEMICAL REACTIONS concentrated solution + solvent dilute solution Based on the number of reagents and products (mainly in organic chemistry) Example: to 50 mL 0.28 M glucose solution enough water given to make a final volume of 1000 mL 50 mL 0.28 M solution 0.050 L × 0.28 molL-1 = 0.014 mol solute final volume: 1.000 L final molarity: 0.014 M initial molarity initial volume Moles of solute = Molarity × Volume = Mi × Vi = Mf Vf final molarity final volume - isomerization: NH4OCN H2NCONH2 - addition: CH2=CH2 + Cl2 CH2ClCH2Cl special cases: dimerization, trimerization, polymerization: 3C2H2 C6H6 - elimination/decomposition: NH4Cl NH3 + HCl - substitution: C2H5Cl +NaOCH3 C2H5OCH3 + NaCl GROUPS OF CHEMICAL REACTIONS Based on the observable change GROUPS OF CHEMICAL REACTIONS element, molecules formed by atoms 2Mg (s) + O2 (g) 2MgO (s) - gas formation: Zn + 2HCl ZnCl2 + H2 - color change: 2KBr + Cl2 2KCl + Br2 element, atom +2 ion -2 ion electron transfer occurs: REDOX REACTION - precipitation: ZnCl2 + H2S ZnS + 2HCl - combustion: CH4 + 2O2 CO2 + 2H2O NaOH (aq) + HCl (aq) NaCl (aq) + H2O (l) +1 ion -1 ion +1 ion -1 ion no electron transfer occurs: ACID-BASE REACTION REDOX REACTIONS reduction – oxidation reaction reduction: the gain of one ore more electrons by a substance – element, compound, ion oxidation: the loss of one ore more electrons by a substance 4Fe (s) + 3O2 (g) 2Fe2O3 (s) oxidation of iron reduction of oxygen EARLIER DEFINITIONS still used widely in organic chemistry reduction: removal of oxygen (O) or addition of hydrogen (H) oxidation: addition of oxygen (O) or removal of hydrogen (H) also reaction with oxygen element (O2) 2Fe2O3 (s) + 3C (s) 4Fe (s) + 3CO2 (g) reduction of iron oxidation of carbon reduction Fe2+ + 2e– Fe Fe3+ + e– Fe2+ Oxidation number (oxidation state) a formal number assigned to every atom in a compound or element, which indicates whether the atom is neutral, electron-rich or electron-poor Fe3+ + 3e– Fe – useful tool for classifying compounds and balancing equations Ag+ + e– Ag – does not necessarily imply actual charges on atoms Cl2 + 2e– 2Cl– O2 + 4e– 2O2– MnO4– + 8H+ + 5e– Mn2+ + 4H2O oxidation ELECTRONEGATIVITY (EN): a dimensionless number showing the ability of an atom in a molecule to attract the shared electrons in a covalent bond F: 4.0 O: 3.5 Cl: 3.0 H: 2.1 Al: 1.5 K: 0.8 (p. 248) Rules for assigning oxidation numbers 1. An atom in its elemental state has an oxidation number of 0. Na, H2, Br2, F2, O2, Ne, S8, P4 oxidation number 0 2. An atom in a monatomic ion has an oxidation number identical to its charge. Na+: +1 Ca2+: +2 Al3+: +3 Cl–: -1 O2–: -2 3. In a heteronuclear covalent bond, the bonding electron pair is assigned to the more electronegative atom. H–Cl H: +1 Cl: -1 H–O–H AlBr3 Al: +3 Br: -1 8. Hydrogen can have an oxidation number of either +1 or -1 depending on the electronegativity of the other element forming the compound. CaH2 Ca: +2 H: -1 NH4+ N: -3 H: +1 H2SO3 H: +1 S: +4 O: -2 5. Fluorine always has an oxidation number of -1 in its compounds. XeF2 Xe: +2 F: -1 SF6 S: +6 F: -1 6. Alkali metals always have an oxidation number of +1, alkaline earth metals +2 in their compounds. LiH Li: +1 H: -1 BaO2 Ba: +2 O: -1 H: +1 O: -2 7. Aluminum usually has an oxidation number of +3. LiAlH4 Li: +1 Al: +3 H: -1 4. The sum of the oxidation numbers is 0 for a neutral compound and is equal to the net charge for a polyatomic ion. H2S H: +1 S: -2 9. Oxygen usually has an oxidation number of -2. XeO2 Xe: +4 O: -2 SO2 S: +4 O: -2 10. Halogens have an oxidation number of -1 in all compounds not containing oxygen or another halogen. FeCl3 Fe: +3 Cl: -1 FeBr2 Fe: +2 Br: -1 Other cases: NaClO2 Na: +1 Cl: +3 O: -2 BrO2 Br: +4 O: -2 HClO H: +1 Cl: +1 O: -2 BrCl Br: +1 Cl: -1 IF5 I: +5 ClO4– Cl: +7 O: -2 ClF3 Cl: +3 F: -1 F: -1 Notable exceptions: oxygen fluorides OF2 O: +2 F: -1 peroxides: H2O2 H: +1 O: -1 superoxides: KO2 BaO2 Ba: +2 O: -1 There are cases when more then one assignment can be used without problems. Fe3O4 Fe: +8/3 O: -2 Fe2: +3 Fe1: +2 O: -2 K: +1 O: -½ REDUCING AGENT: 0 0 +3 • causes reduction -2 4Fe (s) + 3O2 (g) 2Fe2O3 (s) • loses one or more electrons iron: 0 +3 undergoes oxidation REDUCING AGENT • undergoes oxidation oxygen: 0 -2 undergoes reduction OXIDIZING AGENT • oxidation number of atom increases +3 -2 0 0 +4 -2 2Fe2O3 (s) + 3C (s) 4Fe (s) + 3CO2 (g) OXIDIZING AGENT: iron: +3 0 undergoes reduction OXIDIZING AGENT • causes oxidation carbon: 0 +4 undergoes oxidation REDUCING AGENT • gains one or more electrons oxygen: -2 -2 no change • undergoes reduction • oxidation number of atom decreases +1 +7 -2 +1 +4 -2 +1 +6 -2 +1 +7 -2 2 KMnO4 + 5 Na2SO3 + 3 H2SO4 +2 +6 -2 +1 +6 -2 +1 +4 -2 +1 -2 2 KMnO4 + 3 Na2SO3 + 1 H2O +1 +6 -2 +1 -2 +4 -2 2 MnSO4 + 5 Na2SO4 + 1K2SO4 + 3 H2O +1 +6 -2 +1 -2 +1 2 MnO2 + 3Na2SO4 + 2 KOH Mn: +7 +2 decrease, oxidizing agent Mn: +7 +4 decrease, oxidizing agent S in Na2SO3: +4 +6 increase, reducing agent S: +4 +6 increase, reducing agent -5 change 2 mol needed Mn: -3 change 2 mol needed S in Na2SO3: +2 change Mn: 5 mol needed S: +2 change 3 mol needed +1 +7 -2 +1 +4 -2 +1 -2 +1 2 KMnO4 + 1 Na2SO3 + 2 KOH +1 +6 -2 +1 +6 -2 +1 -2 2 K2MnO4 + 1Na2SO4 + 1H2O Fe (s) + Cu2+ (aq) Fe2+ (aq) + Cu (s) Fe: 0 +2 increase, reducing agent Cu: +2 0 decrease, oxidizing agent Mn: +7 +6 decrease, oxidizing agent S: +4 +6 increase, reducing agent Mg (s) + 2H+ (aq) Mg2+ (aq) + H2 (g) Mg: 0 +2 increase, reducing agent H: +1 0 decrease, oxidizing agent Mn: -1 change 2 mol needed S: +2 change 1 mol needed Activity series of elements: strongly reducing more active weakly reducing less active Li K Ca Na Mg Al Zn Cr Fe Co Sn H2 Cu Ag Hg Pt Au react with H2O to produce H2 Au (s) + 3Ag+ (aq) X Au3+ (aq) + 3Ag (s) does not occur strongly reducing more active do not react with H2O react with H+(aq) to produce H2 do not react with H2O or H+(aq) weakly reducing less active Li K Ca Na Mg Al Zn Cr Fe Co Sn H2 Cu Ag Hg Pt Au Fe (s) + Cu2+ (aq) Fe2+ (aq) + Cu (s) Mg (s) + 2H+ (aq) Mg2+ (aq) + H2 (g) Au (s) + 3Ag+ (aq) X Ca (s) + Sn2+ (aq) Ca2+ (aq) + Sn (s) Co (s) + Zn2+ (aq) X SIMPLE APPLICATIONS OF REDOX REACTIONS more active metal + ion of less active metal ion of more active metal + less active metal less active metal + ion of more active metal X more reducing metal + ion of less reducing metal ion of more reducing metal + less reducing metal Combustion: burning of fuel by oxidation with oxygen in air. Fuel: gasoline, natural gas, fuel oil, wood, carbon CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (l) Bleaching: redox reactions to decolorize or lighten colored materials (hair, wood pulp, clothes) Bleaching agent = oxidizing agent less reducing metal + ion of more reducing metal X Batteries: providing electrical current using spontaneous chemical reactions. Redox reactions are needed based on common and inexpensive materials Zn (s) + 2MnO2 (s) + 2NH4Cl (s) usually H2O2, NaOCl, Cl2 Corrosion: deterioration of a metal by oxidation, rusting, usually oxygen from air and moisture is needed 4Fe (s) + 3O2 (g) + 2H2O (l) 2Fe2O3H2O (s) iron rust ZnCl2 (aq) + Mn2O3 (s) + 2NH3 (aq) + H2O (l) Metallurgy: the science of extracting and purifying metals from their ores, always based on redox reactions ZnO (s) + C(s) Zn (s) + CO (g) Respiration: the process of breathing and using oxygen for the many biological redox reactions that provide the energy needed by living organisms. C6H12O6 + 6O2 6CO2 + 6H2O + biological energy glucose (a sugar) Simple acid-base reactions - precipitation - gas formation - neutralization Electrolytic dissociation C12H22O11 (s) H2O C12H22O11 (aq) sucrose molecules no ions in water: NONELECTROLYTE does not conduct electricity - complex formation very important property: electrolytic dissociation NaCl (s) H2O Na+ (aq) + Cl (aq) common salt: ionic solid ions in water: ELECTROLYTE conducts electricity Svante August Arrhenius (1859-1927) HCl (aq) H+ (aq) + Cl (aq) 0.1% Strong electrolytes 'complete dissociation' HCl, HBr, HI STRONG ELECTROLYTE CH3COOH (aq) Weak electrolytes Nonelectrolytes 99.9% CH3COO (aq) + H+ (aq) 99% 1% CH3OH HNO3 HCOOH C2H5OH H2SO4 H3PO4 NaCl WEAK ELECTROLYTE NaOH, KOH precipitation neutralization AgNO3 + NaCl AgCl + NaNO3 HCl + NaOH NaCl + H2O precipitate FeCl3 + 3NaOH Fe(OH)3 + 3NaCl gas formation CH3COOH + NH3 NH4CH3COO 2CsOH + H2SO4 Cs2SO4 + 2H2O complex formation FeS + 2HCl FeCl2 + H2S (g) AgCl + 2NH3 [Ag(NH3)2]Cl Na2CO3 + H2SO4 Na2SO4 + H2O + CO2 (g) HgO + 4KI + H2O K2[HgI4] + 2KOH NH4Cl + NaOH NaCl + H2O + NH3 (g) Zn2+ + 4OH [Zn(OH)4]2 SOLUBILITY: No simple rules for most of the compounds. A compound is probably soluble in water if it contains one of the following cations: Li+, Na+, K+, Rb+, Cs+ NH4+ A compound is probably soluble in water if it contains one of the following anions: Cl, Br, I except that of Ag+, Hg22+, and Pb2+ NO3, ClO4, CH3COO SO42 except BaSO4, Hg2SO4, PbSO4, CaSO4 H2O CH3COOH KBr 'partial dissociation' Pb(NO3)2 + H2S PbS + 2HNO3 HF HClO4 C12H22O11