122ch16d - Assuming the eiementé is me electronggative...

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Unformatted text preview: Assuming the eiementé. is me electronggative than H. as the electronegativy of“. X increases, the X-H bond becomes more po_l_ar and the strength of the a¢i___d_ increases This trend holds true as e lectronggativiy' In______creases a___cross a row of the Qgfiodi c c__hart However. as eiectronegativity decre___a_ses going_ down a fa_mily, acid strength increases because the strength of the H-X bo___nd decreases, even though ' the H-X bond becomes less polar mmilflZflJl w I ,1 ‘ ' 4/ W- am.) ’__I 4 ’ '..A¢z, ' ' ' /1 @ NO; (HMO; lslhe stronger acid because it has more n nprotonated O atoms. so_ I NO,‘ is the stronger base.) [/4/4; 12" b. 4%” of; My. . A A504} (Because P is more electronegative, H3PO4 is the stronger acid. so AsO.1 ' -' is e stron er base. ,4 Jé WW1 W .90 415' ' ‘. ”fir >g/A '1‘ ) #3 3/9 - flange? W 4 based)“ 3 Pa "'4 ,4 fl 3' :I 0032‘ (The ignore ngg‘a ' e theanion, the stron er the attraction for H*.) '— __ g firnge r a 5 e _' 2&2; ; i 4/ 75m)“ [A ’1} X 5190)?!- I‘nc- Ml. I‘M If??? fipX.Wa’ro_¢f) — True. #51514; #"X 6mm; W ' —— False. For onecids with the same structuLe but different central at_o_m. the acid ..._, strength increases as the electronggaylity‘ of the central 59p increases. - *— ——C)— False. HF is a weak acid. weaker than the other hydrogen halides, primarily because '-—— the H-F bond energy is exceptionally high. (d/‘FFJ‘w/f 76 5PM 54/: W) ' _ — i. if ) Theom Acid Base ____ ' Arrhenius forms H+ ions in water produces OH" in water Brensted-Lowry proton (H‘) donor proton acceptor Lewis ' electron pair acceptor electron pair donor The Bransted-Lowry theory is more general than Arrhenius's definition. because it is based on a unified model for the processes responsible for acidic or basic character. and it shows — the relationships between these processes. The ngis umry is M mm still because -—'"' itdoes not restrict the acidic species to compounds having ionizable hydrogen. Any substance — that can be viewed as an electronfiir acceptor is defined as a Lewis acid. — '9’- eZ/omW/wa AW éaJé, __ E Acid Formula Ka(25°C) ~————~———— Acetic Cch00H 1.8 X 10.. w ———- Chloroacetic CHQCICOOH 1.4 x 10—3 " _ —,________ Dichloroacet‘lc CHCIZCOOH 3,3 x 10" Trichlomacetic CCJSCOOH 2 X ‘tO'_1 Using Lewis structures as the basis of your discussion explain the observed trend in acidifies in the series. Ca1- culate the pH of a 0.10 M solution of each acid. The general Lewis structure for these acids is X 80: x—o-o—o——H whereX='HorCI l“. ‘ X . (x —€l Replacement of H by the more. electronegative chlorine atoms causes the central carbon to become more positively charged, thus withdrawing more electrons from the attached COOH' group, in turn causing the 0-H bond to be more polar, so that H” is more readily ionized. |lllll| To calculate pH proceed as usual, except that the full duadralic form must be used for all but acetic acid. . Acid 121-1 acetic 2.87 chloroacetic 1.95 dichloroacefic 1.36 ’g 5311—7 fl trlchloroacetic 1.1 (_.a-——.- [yr ._ _ (1.. /fl " /L1 [’4 .4 ‘I'11r /4141 ’1»- ’ (It any/er M. #4 ’/_ 2) Since H+ does not appear in the overall reaction, it is either a catalyst or an intermediate. An intermediate is produced and then consumed during a reaction, so its contribution to the rate law can usually be written in terms of concentrations of other reactants (Sample Exercise 14.11). A catalyst is present at the beginning and end of a reaction and can appear in the rate law if it participates in the rates. determining step (Exercise 14.64). This reaction is pH dependent because H” is a homogeneous catalyst that participates 'in the rate-determining step. |l|||| ...
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