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5_Reactions.pdf

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Department
Chemistry
Course
CHEM 1E03
Professor
Weisner
Semester
Winter

Description
CHEMISTRY What is the world made of? The material world that we live in is the macroscopic counter part to a microscopic world of atoms & molecules. We will become familiar with the various atoms & the molecules they form. Also, the states of matter & physical or chemical change & the tools we use to describe matter & its changes - measurement & mathematical relationships We begin with a brief review of … __________________________________________________ The Periodic Table Mendeleev (1869) When the elements are ordered according to atomic mass, the chemical & physical properties vary in a periodic fashion eg. Li, Na, K, Rb, Cs 3 11 19 37 55 ← atomic # 8 8 18 18 ← spacing are all similar ⇒ they form a “group”, the alkali metals - a column in the table 1 The position (order) of an element in periodic table = atomic number = # of protons in nucleus of atom KNOW FORM OF TABLE KNOW THE FIRST 20 ELEMENTS → ATOMIC NUMBER, SYMBOL, NAME, & POSITION IN TABLE. Columns form groups → labeled by roman numerals Rows form periods 2 Group I: Alkali metals Li, Na, K, Rb, Cs, Fr soft metals the most reactive metals react with most nonmetals to form ionic salts (not with noble gases though) egs. 2 Li(s) + H (2) → 2 LiH(s) lithium hydride 2 Na(s) + Cl (2) → 2 NaCl(s) sodium chloride - table salt __________________________________________________ Group II: Alkaline earth metals harder, higher melting points (than alkali metals) react more slowly with nonmetals egs. Ca(s) + H (2) → CaH (s) 2 2 Mg(s) + O (2) → 2 MgO(s) Salts of group I and II metals with nonmetals are generally ionic (Be is an exception). → Electrical conductivity 3 Electrical conductivity ionic & covalent solids do not conduct ionic liquids (high T) do conduct covalent liquids do not conduct eg. pure water has an extremely low conductivity However, dissolve a metal halide in H 2 - the resulting solution is a good conductor. ⇒ metal halides are … Electrolytes → produce conducting solutions in water Covalent compounds are Non-electrolytes → produce non-conducting solutions in water Solubility & Precipitation Reactions A precipitation reaction results when an insoluble solid is formed from solvated ions brought together by mixing solutions of soluble substances eg. AgNO 3(aq) + NaCl(aq) → AgCl(s) + NaNO (aq) 3 ↑ solid particles precipitate out of solution 4 Ag (aq) + NO (aq) + Na (aq) + Cl (aq) → − 3 + − AgCl(s) + Na (aq) + NO (aq) 3 ↑ ↑ ↑ precipitates these ions because AgCl(s) remain in solution is an insoluble (in H 2) solid What solids are soluble? Solid classified as … Solubility in water (@ 25° C) soluble ≥ 0.1 mol L −1 −1 sparingly soluble 0.01 to 0.1 mol L insoluble −1 ≤ 0.01 mol L 2− eg. most sulfates (SO 4 ) are soluble exceptions: Ca , Sr , Ba+ 2+ & Pb 2+ sulfates Learn: Table 4.1 on page 139 Solubility refers to equilibrium between solid & ions in saturated solution eg. Ag (aq) + Cl (aq) AgCl(s) net ionic equation corresponding to above rxn Equilibrium lies far to the right – replace by → 5 6 Some Solubility Rules Table 4.1 on page 139 Almost all salts of NH 4 the alkali metal cations are soluble. Otherwise solubility is classified according to anion … Anion SolubleInsoluble ____________________________________________________________________________________ − − NO ,3ClO , 4 CH COO − all 3 (more … eg. HCO 3−& ClO 3− − − − + 2+ 2+ Cl , Br & I all except … Ag , Hg 2 , Cu & 2+ Pb 2− 2+ 2+ 2+ SO 4 all except … Ca , Sr , Ba & Pb2+ _____________________________ 2 + CO ,3 NH 4 &everytheise PO 4−, alkali metal 2− S , cation salts 2− − + O & OH NH ,4alkali metal everything else & larger alkaline ea2+h (beginning with Ca ) cation salts _____________________________________________________________________________________ 7 ACID BASE REACTIONS Arrhenius definition of an acid: produces aqueous solution containing H (aq) + + → actually H3O (aq) is a better description of the H ion in acidic aqueous solution eg. HCl(g) + H O(l) → H O (aq) + Cl (aq) − 2 3 Cl −H O −H → Cl− + H−O −H | | H H Aqueous acid solutions: → sour (eg. vinegar) → change color of indicator (eg. phenolphthalein: red → clear) → react with many metals to produce H (2) + i.e. H3O is an oxidi+ing agent 2+ eg. Zn(s) + 2 H (aq) → Zn (aq) + H (g2 Fe(s) + 3 H (aq) → Fe (aq) + 3/2 H (g) 2 In an acid electrons are pulled away from an H – making the H + more likely to come off as H H − A egs H − Cl, H − S − H δ + δ− δ+ δ − δ+ 2δ− δ+ A is more “electronegative” than H − Contrast: Arrhenius bases form OH (aq) in water eg. NaOH(s) + H O(l)2→ Na (aq) + OH (aq) + H O(l) 2 8 BRONSTED-LOWRY DEFINITION Acid is a proton donor Base is a proton acceptor i.e. an acid-base reaction is a proton transfer reaction + − HA + H O 2 H 3 + A acid base acid base + → an acid has an H which it can donate as H → a base has a lone pair of e s which can accept the proton _____________________________________________________________________________________ Stoncids WA ecids (onyxsamples) _____________________________________________________________________________________ HCl HF HOCl HBr H 2O 3 H2SO 3 HI CH C3OH H2SOH4 3PO 4 O N H 3 2 O l C HO 4 2 H3BO 3 HClO 3 most acids are weak _____________________________________________________________________________________ Learn the strong acids NOTE: strength is not related to concentration 9 Strong acids + − HA(aq) + H O2l) → H 3 (aq) + A (aq) → completely ionized in H O2 → high electrical conductivity Weak acids + − eg. HF(aq) + H O2l) H 3 (aq) + F (aq) → lots of HF in solution → smaller electrical conductivity → weak acids are incompletely ionized i2 H O Redox rxns of acids reduction + 2+ eg. Mg(s) + 2H O 3aq) → Mg (aq) + H 2g) + 2H O2l) oxidation Mg → Mg 2+ + 2e− − + 2e + 2H O3→ H + 2H2O 2 Bases → molecule or ion − → proton acceptor (must have lone e pair) 10 eg. HCl(g) + NH (g) → NH Cl(s) (= NH Cl ) + − 3 4 4 H H − + − δ .. δ δ | .. | :Cl − H + :N − H → :Cl: − + H − N − H .. | | .. H H acid base weak base weak acid In general, HA + B A + HB + Weak Bases: eg. NH (aq) + H O(l) NH (aq) + OH (aq) 3 2 4 H H | .. | .. H − N: + H − O: H − N − H + :O:− | | | | H H H H → incomplete ionization → @ equilibrium - small concentration of NH + & OH − 4 → larger conc. Of NH 3 Strong Bases in H O2 H2O(l) + − eg. NaOH(s) → NaOH(aq) = Na (aq) + OH (aq) − − OH (aq) + H O2l) H 2(l) + OH (aq) 11 eg. NaH(s) + H O(2) → H 2g) + OH (aq) + Na (aq) Na H − + H−O−H → H−H + −O−H + Na + ↑ strong base & reducing agent Strong Bases (ionic) + − Hydroxides: Li OH , NaOH, KOH, RbOH, & CsOH Ca (OH ) , Sr(OH) & Ba(OH) 2 2 2 Not very soluble Mg(OH) 2 is even less soluble however, it dissolves in acidic solution, but notbasic solution => it is a basic hydroxide like the above Corresponding oxides react with H O to form OH− 2 eg. L2 O, CaO Li2O(s) + H O2l) → 2 LiOH(aq) 2− − − O + H−O−H → O−H + O−H + − Hydrides: Li H , NaH, KH, … + − Amides (of group I): Na NH ,2… Aqueous acid-base rxns 12 egs. NaOH(aq) + HCl(aq) → NaCl(aq) + H O(l) 2 KOH(aq) + HI(aq) → KI(aq) + H2O(l) base + acid → salt + water + − + − + − Na + OH + H O3 + Cl → Na + Cl + 2H O2 + − + − + − K + OH + H 3 + I → K + I + 2H 2 − + OH + H 3 → 2H O2 common rxn net ionic rxn Conjugate acid-base pairs:
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