Chemistry Chapter 8: Acids and Bases

Properties of acids and bases

Acid + metal --> H2(g) + salt (ex.2HCl + Mg(s) --> MgCl2 + H2) Acid + metal (hydro-/)oxide --> H2O + salt (these reactions = why acids have corrosive prop.s on most metals, why impt. to keep car battery acid well away from metal car body) Acid + base --> H2O + salt Acid + carbonate --> H2O + salt + CO2 = ALL ARE EXOTHERMIC. Salt = ionic compound formed when H of acid is replaced by metal/othr post. ions, form by reaction of acid + metals/bases, parent acid/base = what reacts to form salt. Net neutralisation reaction: H+(aq) + OH-(aq) --> H2O(l); neutralisation reactions (smtues ysefyk ti reduce effect of acid/base, ex. treating acid indigestion using "antacids" of mix w. weak alkalis ex. Mg(OH2) nd Al(OH)3) w. bases = exothermic; enthalpy of neutralisations = enthalpy change when acid + base react to form 1 MOL WTR; for reactions betweel all strong acids nd bases ΔH = very similar to -57 kJ/mol, approx.ly, bc. all have same net reaction involving formation of wtr from its ions nd nearly all go through full reaction. Large range of metal reactivity when adding small metal piece to dilute sol. of acid (ex. top Gr. 1 like Na, K react too violently, less reactive metals like Au/Ag usually don't react at all = partly why are so valualbe bc. more corrosion-resistant). Hwvr, common nitric acid HNO3 reacts w. metals but usually w.o. releasing H2 due to oxidising prop.s Alkalis = slippery, remove fats/oils from fabrics, when dissolved in wtr release OH- Spectator ions = NOT involved in reaction, to find spectator ions 1st dissociate compounds: HCl + NaHCO3 --> CO2 + H2O + NaCl = H+ + Cl- + Na+ + HCO3 --> CO2 + H2O + Na+ + Cl- (bc. occurs in wtr so some are dissociated), then cross out species occurring twice --> H+ + HCO3 --> H2O + CO2 (everythin aq.) = what reaction actually happens

Acid-base titrations

Acid-base neutralisation reactions can be investigated quantitatively using titration (=one of most widely used procedures in chem. quality control of food/ drink, health/ safety checks in cosmetic industry, clincal analysis in med.). When making titration curve has burette + pH meter --> curve w. pH on y-axis, vol. of added acid /base on x-axis. Ex. vinegar of unknown conc., vol carefully measured w. pipette, add small amount of ex. NaOH vol.s, read off how much added + resulting pH --> plot values on graph. Strong acid titrated agains strong base has very straight line from ca pH 1 until all of a sudden is asymptote straight up until pH very high up, pH jump. ex. from 3-11 and pH = equivalence p.; strong acid + weak base = looks same at first but when reaches pH<7 during jump it flattens out remarkably (flat slope bc. weak base so has strong conj. acid, ex. NH4+, which will keep pH down when reacts w. wtr into NH3 + H2O which will keep pH down as more base is added; equivalence p. uncertain in these curves so recommended to use strong acid/ base when doing titrations to find anthr acid's/base's conc.); weak acid + stron base = starts higher up, around 4 or similar, nd has steeper curve up to equivalence p. (has so bc. OH- form from the strong bases's conj., which makes the pH incr.) when ahs sharp rise until flattens out at top again (jump ex. from 7-11); weak acid + weak base = starts where weak acid + strong base began but at equivalence p. continues out w. strong acid + weak base Equivalence point = stoichiometric point when sol.s exactly neutralise When base added to acid in neutralisation is change in pH, but change doesn't show linear relationship w. vole. of base added, partly bc. pH scale is logarithmic; record pH w. pH meter/ data-logging device as function of vol. of base added, plot as pH curve; pH curves can also be derived from theory by calc.in pH at diff. vol.s of base added, but more laborious process but w. same results. In most titrations big pH jump occurs at equivalence p. = p. of inflection, is always half-way up this jump Strong + strong base --> pH at equivalence = 7 as neither ions hydrolyse enough as are fully dissociated strong acids/bases; as base is added to acid neutralisation of some of acid occurs while excess acid remains until equivalence where amounts of acid nd base fully react; after that mixture contains excess base, as vol. changes during addition this must be taken into account when determining conc.s. Small addition of base around equivalence p --> dramatic rise in pH; initial pH can be found from y-intercept, hence also [H3O+]; after eq. p. curve flattens out at high pH value (of strong base) Weak acid + strong base --> pH at eq. p.>7 as anion hydrolysis releases OH-; initial pH = fairly high, stays rel.ly constant until eq. is reached (=buffer region); is jump in pH at eq. from pH 7-11 which is smaller jump than in strong-strong, after eq. curve flattens out at high value; half-equivalence p. = when exactly half of first sol. has reacted nd neutralised - mixture has equal quantities of weak acid nd its salt form w. base = buffer sol. = why pH in this region is shown to be rel.ly constant to change in pH on addition of small amounts of base Calc.ing pKa from half-equivalence p. = pH = pKa or pOH = pKb (pKb can be calc.d when titration done w. acid added to base) Strong acid + weak base --> pH at eq. p.<7, pH stays rel.ly constant through buffer region to equivalence jump in pH = smaller, after eq. p. curve flattens out at fairly low pH Weak acid + weak base --> pH at eq. p. more difficult to define bc. is rel.ly flat slope overall: initial pH fairly high, adding base --> steady rise in pH, change in pH at eq. p. = MUCH LESS sharp than in othr titrations; after eq. p. curve flattens out at fairly low pH (of weak base). Due to sev. eqm being involved titration doesn't clearly define eq. p. as is no sig. pH jump --> better to use othr tech.s, ex. conductimetricm measurements to determine eq. p. If pH curve desrcibes addition of acid to base curve will be inverse, starting high nd ending low, but same things apply For calc.s, assume full dissociation: 1. calc. n1(acid initial) = vV, then n2(base added) = c2V2; 2. then n(acid remaining after base addition) = n1 - n2 = [n(H3O+)]; 3. then n(H3O+)/V(new vol. of acid + base) = [H3O+]; 4. repeat this for diff. vol.s of base added (See book for clarification!!!).

Acid deposition

All rain wtr nat.ly acidic bc. dissolves CO2as H2O (l) + CO2(g) <--> H2CO3(aq), which ionises to form eqm H2CO3 <--> H3O+ + HCO3- --> sol. of at least pH 5,6; acid rain = sol.s w. pH BELOW 5,6 so therefore contains additional acids. Main contributors to acid rain = oxides of S nd N (= primary pollutants. Acid rain = secondary pollutant prod.d when these acidic gases dissolve in wtr (pariculates act as catalysts in prod. of secondary pollutants) Acid deposition = broader term than acid rain, incl.s all processes by which acid components as precipitates/gases leave atmosphere; 2 main types of acid deposition = wet acid deposition (rain, snow, sleet snöblandat regn, fog, mist, dew; fall to ground as aq. precipitates), dry acid deposition (acidifying particles, gases, gall to ground as dust/ smoke to later dissolve in wtr to form acids, typcally occurs near source of emission, wet acid deposition = dispersed over much larger areas/distances from emission source; formed in air from SO2 nd NO emitted by thermal pwr stations, industry, motor vehicles; major source = burning fossil fuels; pollutants carried by prevailing windsm converted/ oxidised into H2SO4 nd HNO3 to be dissolved in cloud droplets to fall to ground as acids) SO2 from burning fossil fuels (espec. coal nd heavy oil in per plas for electricity +released in industrial processes of smelting where metals extracted from ores), estimated 50% of annual SO2 emissions from coal burning. S(s) + O2(g) --> SO2(g) = colourless gas, sharp/ putrid smell SO2(g) + H2O(l) --> H2SO3(aq) = sulphurous acid SO2 can also oxidise into SO3: 2SO2(g) + O2(g) --> 2SO3(g), then reacts w. wtr: H2O(l) + SO3(g) --> H2SO4(aq) Sev. mechanisms might occur in these reactions w. complex chem. given wide range of conditions/ othr chem.s in atmosphere; during sunlight h PHOTOOXIDATION may occur (oxidation may be catalysed by tiny metal particles ex. Fe/Mn present in clouds). O3 or H2O3 as pollutants in atmosphere can be involved; more detailed study incl. short-lived hydroxy free radicals (HO•) from reactions between wtr nd O• or O3, ex. OH• + SO2 --> •HOSO2, then •HOSO2 + O2 --> •HO2 + SO3 NO = mainly prod.d from internal combustion engines (burning fuel --> heat E --> N2(g) + O2(g) --> 2NO(g), ΔH = +181 kJ/mol). Similar reaction directly creates brown gas NO2: N2(g) + 2O2(g) --> 2NO2(g). NO2 also forms from oxidation of NO: 2NO(g) + O2 --> 2NO2(g); NO2 dissolves in wtr to form mix of nitrous acid HNO2 nd nitric acid HNO3: H2O(l) + 2NO2(g) --> HNO2 (aq) + HNO3(aq) OR NO2 can be oxidised to form HNO3: H2O(l) + 4NO2(g) + O2 --> 4HNO3; these to reactions = prod.d by diff. mechanisms, dep. on conditions/ chem.s in atmosphere - photo oxidation, ozone presence, •HO all contribute to prod. of nitrous/ nitric acid (•HO + NO --> HNO2, •HO + NO2 --> HNO3) --> acid rain (main active components is thus = H2SO3, H2SO4, HNO2, HNO3 (aq)) Building mat.s marble nd limestone = both forms of CaCO3; both SO2 in dry deposition AND H2SO4 in acid rain react w. this into CaSO4 (then CaCO3 + H2SO4 --> CaSO4 + H2O(l) + CO2(g) OR 2SO2 + O2 + 2CaCO3 --> CaSO4 (aq)) Bc. CaSO4 = somewhat more soluble than CaCO3 it washes out of limestone/flakes off marble; CaSO4 = GREATER molar vol. than CaCO3 --> its formation causes expansion nd stress in stonewrk; similar reaction occurs w. nitric acid --> soluble nitrate salt: CaCO3(s) + 2HNO3(aq) --> Ca(NO3)2 + H2O(l) + CO2 --> these reactions --> erosion of structures --> loss of ex. many hist. buildings/statues. Acid deposition --> corrosion of metals; both dry deposition + acid rain react w. metals, ex. Fe -- salt --> enables ionic conductivity --> incr. in rate of electrochem.corrosion reactions like rusting: Fe(s) + SO2(g) + O2(g) --> FeSO4(s) OR Fe(s) + H2SO4(aq) --> FeSO4(aq) + H2(g) Also, acid rain can react w. (=remove) protective oxide layer on surface of metals like Al: Al2O3(s) + 6HNO3(aq) --> 2Al(NO3)3(aq) + H2O(l) (this reactant releases MORE of pollutant NO inst. to form MORE acid rain, bc. HNO3 doesn't react as typical acid reacting w. metal into releasing H2 bc. is oxidising agent) ACid rain --> slower plant growth + injury/ death of plants; ex. causes impt. imperals ex. Mg2+, Ca2+, N+ held in soil to become soluble --> wash away in process called leaching before can be absorbed by plants. W.o. sufficient Mg2+ plants can't synthesise chlorophyll --> can't photosynthesise sufficiently. Acid rain also --> release of substances toxic to plants, ex. Al3+ damaging plant roots => acid rain --> sig. damage to metallic structures like bridges, rail road tracks, vehicles Dry deposition can directly affect plants by blocking stomata for gas exchange. Forests in hilly regions appear espec. vulnerable bc. tend to be surrounded by acidic clouds nd mists -sme of worst effects on forests been in Europe, but impact seen globally Acid rain --> sev. lakes unable to supp. life, ex. trout, perch, mny othr fish caN'T survive at pH below 5; below H rivers = effectively dead as toxic Al3+ normally trapped in rocks as insoluble Al(OH)3 leach out under acidic conditions (Al(OH)3 + 3H3O+ --> Al3+ + 3H2O); Al3+ interferes w. operation of fish gills nd reduces ability to take in O2; acid rain contributes to additional prob. eutrophication (=over-fertilisation of bodies of wtr, can be caused by nitrates present in rain --> algal blooms --> O2 depletion --> smtimes death of lake/stream) Acid rain doesN'T directly affect hum. healty but components can react to form fine sulphate nd nitrate particles able to travel long distances to be present in inhaled air, can irritate respiratory tract, incr. risk of illnesses like asthma, brocnhitis, emphysema + cause irritaion to eyes. Debase of toxif metals ions like Al3+, Pb2+, Cu2+ by reaction of acid rain on metal structures like pipes = also potential health risk. Link between industrialisation --> atmospheric pollution nd acid rain = described in Manchester 1852, but phenomenon w. little attention until 1970s when impacts evident in sev. countries --> govt.al measures to reduce S/NOx emissions Reduction of SO2 emissions: Pre-combustion methods: reduce/remove S present in coal/oil before combustion; where S is present as metal sulfide can be removed by crushing coal nd washing w. wtr; high density metal sulfide sinks to bottom so sep.s from clean coal. Hydrodesulfurisation HDS = catalytic process removing S from refined petroleum prod.s by reacting it w. H2 to form H2S = highly toxic gas so is captured nd later converted into elemental S for manufactureing H2SO4. Post-combustion methods: flue-gas desulfurisation can remove up to 90% of SO2 from flue gas in smoke stacks of coal-fired pwr stations before is released into atmosphere; process uses wet slurry of CaO nd CaCO3 reacting w. SO2 into neutral prod. CaSO4: CaO(s) + SO2(g) --> CaSO3(g) CaCO3(g) + 3O2(g) --> CaSO3(s) + CO2(g) 2CaSO3(s) + O2(g) --> 2CaSO4(s) (prod. has industrial uses as making plasterboard) Reduction of NOx emissions: Catalytic converters in vehicles: exhaust gases can be controlled by use of catalytic converters where hot gases mix w. air nd passed over Pt/Pd- based catalyst, reaction converts toxic emissions --> rel.ly "harmless" prod.s, ex: 2CO(g) + 2NO(g) --> 2CO2(g) + N2(g) Lower-temp. combustion: formation of NO reduced at lower temp.; recirculating exhaust gases back into engine lowers temp. to reduce NOx in emissions. Othr options: in addition to finding ways to reduce primary pollutant emission when fossil fuels burn, othr sol.s to prob.s of acid deposition = lowering demand for fossil fuels => more efficifent transfer syst.s, greater use of public transport, switching to renewable E. Restoring damaged ecosyst.s = long-term process, ex. use CaO(lime) or Ca(OH)2 to neutralise acid in wtr

First definition of acids

Arrhenius 1887 - acid(aq) --> anion(aq) + H+(aq); base(aq) --> OH-(aq) + cation(aq); suggested acid/base dissociates in twr (close to theory used today to xplain acid/bases, but his focus was only in aq. sol.s so req.s broader theory Acid-base theory = ex. central to wtr/air pollution, global warming's effects on ocean chem., drug action in body etc.; acid encounters = ex. wine smell exposed to air, sour grapefruit taste, acid indigestion; alkailis/bases = baking soda, household cleaners w. ammonia, med. against indigestion. Most acid-base reactions incl. eqm In acid-base theory ionisation nd dissociation = typically interchangeable ALL nitrates = soluble Na-, K-, NH4+- compounds = generally soluble ALL sulphates = soluble EXCEPT for BaSO4, PbSO4, CaSO4 ALL chlorides = soluble EXCEPT AgCl, PbCl ALL carbonates = INSOLUBLE EXCEPT Na-, K-, NH4+ ONLY SOLUBLE hydroxides = NH4OH, LiOH, KOH

Lewis acid-base theory

Broader than B-L theory (this ofc. often incl.s H+ transfer due to its vacant orbital, but also molecules w. incomplete valence shells --> always lead to formation of cov. bond as coordinate bnd bc. both e-s come from base --> ligands = Lewis bases donating e- pair to transition metals; typical ligands in complex ions = H2O, CN-, NH3 bc. possess lone e- pairs); Lewis realised that to be base must have lone pair of e-, theory focuses on e- pair inst. of H+ --> Lewis BASE = e- pair DONOR (=nucleophile, attracted to pos., used lots in organic chem., e- rich); Lewis ACID = e- pair acceptor (=electrophile, per definition attracted to neg., e- deficient), term usually reserved for species that ONLY can be described using Lewis theory (acids that DON'T RELEASE H+), incl wide range of in-/organic species not recognised in B-L theory, but in many cases B-L is sufficient (=more useful for describing acid-base reactions, espec. in aq. sol.s). Lewis base = B-L base, ex. w. Lewis structures: NH3 (N w. 1 lone pair) + H2O(w. 2 lone pairs), NH3's lone pair is transferred to H on H2O in joint compound w. coordinate bond as NH4. This e- transfer = indicated by curly arrow, show where lone e- pair goes, starts at lone pair, arrow head = where pair ends up; bc. the H on H2O can't bond to H2O and NH4 at same time is othr curly arrow from start of line signalling its bond to O and arrow head goes to the O (small arrow) to show that new prod. OH- receives one of H's protons - in momemnt of bonding H's bond to H2O vanishes so H gives up this cov. bond of 2 e- to O as lone pair of e-. Mostly, Lewis nd B-L acids overlap bc. usually incl.s H+ transfer, is Venn diagram w. bigger circle = Lewis acids nd within this is smaller circle w. B-L acids; B-L bases = ALWAYS Lewis bases. Acids = huge group, some are both Lewis nd B-L acids, some are ONLY LEWIS ACIDS - ex. transition metals in ligands or species w. incomplete octet. Ex. BF3 = incomplete octet w. only 6 valence e- --> eager to accept lone pair of e- = e- pair acceptor --> why ex. easily bonds w. NH4+ for lone e- pair in coordinate bond (shown w. arrow TOWARDS acceptor). Complex ions forming acids = also Lewis acids, ex. Cu2+(aq) + 6H2O(l) --> [Cu(H2O)6]2+ (aq) - Cu accepts 6 lone e- pairs from H2O.

Amphiprotic substances

Can act as BOTH as and base, donate nd accept H+; amphoteric= can REACT w. both acids and bases. All amphiprotic species = amphoteric, but all amphoteric species aren't necessarily amphoprotic (ex. Al2O3 = amphoteric but NOT amphiprotic). Amphiprotic espec. relates to B-L acid-base theory where emphasis on proton transfer. Amphiprotic = H2O (when acts as base its conjugate acid = H3O+; when acts as acid its conjugate base = OH-); HCO3- (when acts as base its conjugate acid = H2CO3, when acts as acid its conjugate base = CO3 2-); HSO4- (when acts as base its conjugate acid = H2SO4, when acts as acid its conjugate base = SO4 2-)

pH

Can only compare 2 acids/bases if are in same conc. bc. pH measures conc. of H+ so if ex. CH3COOH(aq) is in much higher conc. than HCl it will have more mols per dm3 nd hence also more H+s dissociated; pH = inversely related to [H3O+], 1 pH incr. = DECR. in [H3O+] by 10x (sig. fig.s in logarithm = only ouny fig.s to RIGHT of decimal p.) pH = -log[H3O+] --> [H3O+] = 10^-pH. Bc. pH scale = mathematical concept goes infinitely in both directions, but 1-14 = most common pH range High pH = low [H3O+], low pOH, inst. = high [OH-]; low pH = high [H3O+], high pOH, low [OH-]; scale is logarithmic. Above/below scale goes pOH in opp. direction of pH scale (so 14 on pH scale = 0 on pOH, 0 on pH = 14 on pOH. Using scale, if ex. pOH = 12 its [OH-] must be 10^-12 bc. pOH = 10^-pOH, at this p. pH must be 2 (bc. pOH + pH = ALWAYS 14 at 25*C), so [H3O+] must be 10^-2 bc. [OH-] _ 10^-pH, if multiplied = 10^-14. By measuring the pH and plugging into formula we can get [H3O+] and compare this to the initial molar conc. of the sol. (ex. CH3COOH) and thereby see how many ions have dissociated (for 0.1mol/dm3 CH3COOH will be 10^-3 = 1%) At neutral, [H3O+] = [OH-] (why neutral circulates around 7 at normal temp.s)

Ionisation of water

Dissociation of wtr = endothermic, bond breaking, so incr. in temp. --> incr.d Kw value --> incr.d [H3O+] nd [OH-] conc.s so lowered pH. Possible to have ex. low [H3O+] in a strong base bc. of this. Wtr = amphiprotic tgthr w. acid/base acting as acid/base. --> ex. in pure wtr w. 2 wtr molecules they will behave as such as to react w. each othr amphiprotically bc. can't react w. anything else --> H2O(l) + H2O(l) <---> (EXTEMERLY to left) H3O+(aq) + OH-(aq) acting as base nd acid respectively = happens to EXTREMELY SMALL EXTENT, but still occurs nat.ly at an EQM, very much to left so we can actually consider wtr conc. to be constant bc. reached so quickly. Bc. is eqm we can calc. Kc: Kc. = [H3O+]eq[OH-]eq/[H2O]eq^2 Bc. we consider [H2O] to be constant we multiply to get: Kc[H2O]eq^2 = [H3O+]eq[OH-]eq = new constant as wtr conc. basically is constant --> Kw = [H3O+][OH-] = IONIC PRODUCT OF WTR = Kw (or [H+][OH-] = Kw) Under STANDARD CONDITIONS (298K, 25*C), Kw =1,00 x 10^-14 mol2dm^-6. By knowing Kw nd conc.s we can calc. pH of wtr at 25*. Bc.[H3O+] nd [OH-] are equal each othr at neutral we can rewrite ionic prod. of wtr as Kw = [H3O+]^2 --> √Kw = [H3O+] = √(1,00 x 10^-14) = 1.00 x 10^-7 pH = -log[H3O+] = -log(1.00 x 10^-7) = 7,00 = why wtr is neutral at 7 in 25*C. BUT bc. this is eqm expression varies w. temp., but it can still be neutral even if is 6,89 if that means [H3O+] = [OH-] H2O(l) + H2O(l) + E (through incr.d temp.) <--> H3O+(aq) + OH-(aq), ΔH = +57,3 kJ/mol = ENDOthermic --> so if we incr. temp. we shift eqm to right by creating more H3O+ nd OH-, resulting in a slightly lower pH as Kw thereby incr.s. By taking -log in ionisation of wtr expression: -logKw = -log[H3O+] + -log[OH-] Bc. -log[H3O+] = pH we get --> -logKw = pH + -log[OH-] --> pKw = pH + pOH Bc. Kw = 1,00 x 10^-14 at 25*c, pKw (-logKw) = -log(1,00 x 10^-14) = 14 --> Kw = 14,00 = pH + pOH pH = -log[H3O+] --> all are dissociated in in strong acid/base --> [H3O+] = c(acid) (in ex. HCl + H2O --> H3O+ + Cl- H3O+:HCl in 1:1 molar ratio); same thing for base - strong base [OH-] = c(base); if c = 0.1 mol/dm3 --> pOH = - log(0,1) = 1; pH + pOH = 14 --> PH = 14 - 1 = 13 For weak acids is more complicated to calc. pH: For acid HA + H2O <--> H3O+ + A- Kc = [H3O+]eq[A-]eq/[H2O]eq[HA]eq --> Kc[H2O]eq = [H3O+][A]eq/[HA]eq --> Acid dissociation constant/acid ionisation constant Ka bc. [H] is basically constant (eq can be omitted bc. eqm is reached so fast now) Ka = [A-][H3O+]/[HA] Every acid has own Ka value (higher value = stronger acid, VERY WIDE range of values - to list smaller range of values in lists of Ka values pKa is typically expressed by, but pKa = -logKa --> Ka = 10^-pKa), Ka = fixed fo particular acid at specified temp., value dep.s on position of eqm of acid dissociation so gives direct measure of acid sterngth, doesn't dep. on acid conc. not presence of othr ionschanges w. temp. bc. is eqm constant; lower pKa value = stronger acid bc. thereby higher Ka value as pKa is logarithmic and that means is higher [HA] conc. so fewer are dissociated; Strong acids don't have Ka values bc. are completely dissociated so denominator [HA] = 0. To cal. pH of weak acid, find Ka value from table + conc.; ex. 0.1 mol/dm3 CH3COOH Ka ~ 1,74 x 10^-5 In conc. tables for eqm expression c(in) for CH3COOH = 0, for H2O = excess, for H3O+ nd CH3COO- = 0; for all Δc = -/+ x; c(eq) CH3OOH = 0,1-x nd x for dissociated ions Bc. Ka is extremely small, outside our sig. fig. range, we will get x to be extremely small as well --> we assume 0,1-x~0,1 (ALWAYS STATE assumptions made nd justify why) --> [CH3COOH]eq = [CH3COOH]in since Ka<10^-3 ---> Ka = [H3O+][CH3COO-]/[CH3COOH] --> 1,74 x 10^-5 = x^2/0,1 --> x ~1,32 x 10^-3 mol/dm3 --> [H3O+] = 1,32 x 10^-3 --> pH = -log(1,32 x 10^-3) = 2,88 Exact same applies for bases but for Kb base dissociation constant, fixed value for particular base at specificed temp., gives measure of base strength, relates to eqm position (--> base strength); higher Kb value at particular temp. = greater ionisation --> stronger base. Kb incr.s = strong base, pKb decr.s = strong base; 1 unit change of pKa nd pKb = 10x incr. in Ka or Kb value Ka x Kb = [H3O+][A-]/[HA] x [HA][OH-]/[A-] --> [H3O+] x [OH-] = Kw --> Ka x Kb = Kw --> -log on both sides --> -logKw = -logKa + -logKb At 298K Kw = 1,00 x 10^-14 --> pKw = 14,00 --> pKa + pKb = 14,00 at 298K = holds for any conjugate acid-base pair in aq. sol.s , so from Ka value we can calc. Kb for its conjugate base --> stronger acids have weaker conjugate bases nd vice versa; weaker acid = further eqm to left so stronger conjugate base Greater Kb = more ions = stronger base = smaller pKb In some tables only pKa available, for bases this is tha pKa of their conjugate acid - ex. for NH3 pKa is for NH4+, but we can find pKb by taking 14 - pKa

Conjugate acid-base pairs

In all acid-base reactions are conjugate acid-base pairs - in ex. (aq) HCl + H2O --> Cl- + H3O+, HCl and Cl- are conjugate acid-base pairs, and H2O + H3O+; always differ by 1H+ = typical P1 q. to identify conjugate pair (conjugate bases have 1 LESS H+ than its conjugate acid; conjugate acids have 1 MORE H+ than conjugate base!!!). HA + B <--> A- + BH+, in forward reaction HA = acid bc. donates H+, B = base, but in reverse reaction is vice versa bc. BH+ donates H+ to A nd A- receives H+ like base Conjugate acid-bases = ex. if you reverse above reaction Cl- would act as an acid in relation to HCl bc. Cl- accepts the H+ to become HCl and the H3O+ is the acid bc. it donates the H+.

Buffer solution

Reduces impact of one thing on anthr, "shock absorber", in acid-base chem. reduces pH impact of added acid/base on chem. syst. --> buffer sol. resistant to changes in pH on addition of SMALL AMOUNTS of strong acid/alkali (ofc if add infinite acid/base pH will change). Wtr = very vulnerable to sig. pH fluctuations upon adding ex. only 0.1 cm3 1 mol/dm3 HCl --> major impacts on cham. reactions in aq. sol.s - bio. syst.s can operate efficiently only within narrow pH range, basically bc. pH's effect on enzymes/ biochem. reactions --> these syt.s dep. on buffers; mammalian blood = excellent complex nat. buffer/ ocean maintaining suitable conditions for life (pH 7-8 by sev, buffer syst.s), ex. electrophoresis, fermentation, dyes industry, instrument calibration dep. on buffers Buffers maintain pH close to sey value, all don't wrk at pH 7, diff. buffer sol.s can be made at nearly any pH; when make buffers you can choose what to mix to maintain certain pH 2 main types of buffers: acidic (retain pH<7; made by mixing aq. sol of weak acid + weak sol. of its salt w. strong alkali, ex. CH3COOH + NaCH3COO, which is its conjugate base when it reacts w. NaOH) nd basic (retain pH>7) buffers (+SOME retain pH 7), both = ,ixtures of 2 sol.s composed of 2 species of conjugate acid-base pair; consist of weak acid AND weak base, usually conj. acid-base pairs, typically in high conc.s nd similar amounts. In above acidic buffer (aq), ex.: NaCH3COO --> Na+ + CH3COO- to 100% CH3COOH + H2O <--> CH3COO- + H3O+ to 1% --> if we mix same amount of salt nd acid we'll have abt. equally many ethanoate ions as ethanoic acid (100% vs. 99%) --> lots of ethanoate ions shift the eqm reaction slightly left so they are abt. equal in amount --> [CH3COOH] ~ [CH3COO-] in a buffer sol. --> is conj. acid-base pair in sol. If add strong acid we add lots of H3O+; in buffer sol. conj. base CH3COO-, reacts w. H3O+ from strong acid into CH3COO- + H3O+ <--> CH3COOH + H2O, so we just create more CH3COOH which per se can't change the pH w.o. having smthing to act as a base, and if it reacts w. the wtr again the eqm will shift right again to compensate for that. Mixture consists rel.ly high conc.s of both CH3COOH nd CH3COO-, acid + conj. base, can be considered reservoirs ready to react w. most of added OH- nd H3O+ respectively in neutralisation reactions (OK to have free ex. CH3COO- in sol., if don't have H3O+ won't make diff.) If add strong base to the buffer sol. the acid of conj. acid-base pair takes care of strong base: ex. CH3COOH + OH- <--> CH3COO- + H2O ---> pH largely stays same as added H3O+ nd OH- used up in these reactions they don't persist in sol. Basic buffer sol., ex: NH3 + H2O <--> NH4+ + OH- (1%) NH4Cl (of weak base NH3 + strong acid HCl) --> NH4+ + Cl- (100%) --> [NH3] ~ [NH4+] --> as we add NH4+ we add more of prod. so shift eqm left so are same conc.s of both --> [NH3] = [NH4+] (conj acid = conj. base) If we add acid, H3O+, will react w. conj. base NH3 into H2O + NH4+ (which won't change pH further than before), if add base, OH-, will react w. NH4+ into NH3 + H2O; mixture = rel.ly high conc.s of NH3 nd NH4+ (base + conj. acid) --> act as reservoirs to react w. added H3O+ nd OH- respectively in neutralisation reactions Most often when making buffer sol.s we want it to have pH close to no. --> start w. weak acid/base w. pKa as close as possible to desired pH (if has pKb value find its pKa of its conj. acid, must be pKa for pH measures, can't be pKb - in making buffer sol.s can use pKb if it inst. is close to the desired pOH of buffer); when deciding components of buffer take weak acid/base + salt w. its conj. in ca same molar amounts, mix them in same beaker; if have exact equal no. of moles (n(acid) = n(base)) then pH of sol. = pKa of sol. bc. cancel out (see eq. below), if have little more of acid pH will be slightly lower, if more of base pH slightly higher Can also create buffer by starting w. only weak acid/base, then perform neutralisation through titration w. strong base/acid, but only for little while so ca half of it been, neutralised; ex. CH3COOH(aq) + NaOH(aq) --> NaCH3COO- (we are prod.ing its conj.) If we have 2 mol of acid first, we add 1 mol of NaOH bc. then that results in n(buffer) having 1 mol CH3COOH, 0 NaOH, 1 NaCH3OO + excess H2O --> start w. twice as much of weak acid as strong base pH of buffer dep.s on interactions of its components - pKa or pKb of its acid/base, ratio of initial conc.s of acid/base nd salt Dilution doesn't change pH of buffer but lowers its buffering capacity - kA nd Kb = eqm constants so areN'T changed by dilution, which doesn't change ratio of acid/base to salt conc. either as both components will be decr.d by same amounts --> diluting buffer doesN'T change H, but alters AMOUNT of acid/base it can absorb w.o. sig. changes in pH (buffering capacity) which dep.s on molar conc.s of its components so decr.s as they are lowered by dilution As temp. affects Ka/b value also affects pH of buffer = why constant temp. should be maintained in all wrk w. buffers, as calibration w. pH meters. Temp. fluctuations must also be minimised in many med. procedures ex. blood transfusions bc. effect on buffers in blood. Through pKa/b values becomes clear weaker acid --> stronger conj. base

Strong vs. weak acids

Strength of acid/base = measure how reaily dissociates in aq. sol.s = inherent prop. of particular acid/base, dep.s on its bonding. Strong acids = high conductivity bc. nearly 100% of molecular compounds dissociate (ex. HCl(aq)), but weak acid ex. CH3COOH only ca 1% dissociate into H3O+ + CH3COO- (bc. latter NOT much more stable than parent acid; has eqm very much to left; Cl- = very stable as ion). Acidic sol = contain H3O+ (hydrooxonium/ hydronium/oxonium) ions; Diprotic = donate 2 H+s, ex. H2SO4; monoprotic= donate 1 H+, ex. HCl, HNO3; these 3 = strong acids, rest are weak acids (only slightly dissociate into ions, ex CH3COOH, H2CO3); these are molecular compounds Strong bases = nearly all split into ions (are already ionic), ex. NaOH; all Gr. 1 hydroxides = strong bases (espec. Li-, Na-, K- etc. + Ba(OH)2) Weak bases = all MOLECULAR, only slightly dissociate, ex. NH3 + H2O <--> NH4+ + OH-, ETHYLAMINE C2H5NH2 (such amines can be considered organic derivatives of NH3 in which one of H atoms has been replaced by alkyl hydrocarbon gr.) or Fe(OH3) (on test 1. state w. words <--> implies mostly go reverse). If we have very strong acid its conjugate base (ex. Cl- for HCl= bascially useless as conjugate base (ex. Cl- for HCl) bc. so weak in compaison that reaction can't go backwards; for ex. CH3COOH its conjugate base = equally nd opp.ly weak as an acid so it continues in driving reaction backwards (bc. can't accept that many H+s). Base ionisation reactions favour prod. of weaker acid conjugate Strong acids = faster rate of reaction, seen ex. in how fast Mg dissolves and bubbles create (H2 gas). Ex. 2HCl + Mg(s) --> H2(g) + MgCl2 (aq) is actually: Mg(s) + 2H+ + Cl- --> Mg2+ + H2 + Cl- (like displacement reaction, then Mg2+ reacts w. Cl- as wtr evaporates). Rate of reaction dep.s on H+ conc. --> HCl nearly dissociates entirely into H+ + Cl- so [H+] is much higher than that of CH3COOH's so reacts faster bc. the H+ is what reacts, but it we let both samples go to completion both will do so. For CH3COOH as it reacts w. Mg will continually shift eqm right bc. as the H+ are used up on the right new ones are prod.d (to keep ca 99:1 ratio) acc. to Le Chatelier's principle (=why all Mg eventually is used up) Rate of reaction dep.s on H+ conc.s --> HCl(aq) + H2O(l) --> H3O+(aq) + Cl-(aq) = more correct but equally accepted as HCl(aq) --> Cl- + H+(aq)

Neutralisation reaction/ salt hydrolysis

acid + base --> salt + wtr HA + MOH --> M+A- (ionic compound) + H2O Does, hwvr, not result in neutral sol., dep.s on salt created nd if its ions react w. wtr Salt hydrolysis can occur to change pH of sol., dep.s on to which extent its conj. acids/bases react w. wtr nd hydrolyse it to release H3O+ or OH-. Combo of strong/weak acid + strong/weak base decides resulting acidity of sol. Strong acid + strong base at 298K: HCl(aq) + NaOH --> NaCl + H2O --> neutral sol. bc. NaCl dissociates into Na+ + Cl- which doN'T react w. wtr bc. are conj. acids/bases to strong acids/bases so are themselves very weak so can't react to go back to parent acid/base --> strong + strong = neutral (Na+ NOR Cl- --> NO HYDROLYSIS) Strong acid + weak base at 298K: HCl(aq) + NH3 <--> NH4Cl + H2O NH4Cl --> NH4+ + Cl- => Cl- obv. doesn't react w. wtr, but resulting NH4+ does --> salt hydrolysis occurs bc. is a conj. acid to weak base so is strong own conj. acid on own --> NH4+ + H2O <--> NH3(aq) + H3O+ --> resulting sol. = acidic ALSO (MUST KNOW THIS EX.!!!): HCl(aq) + Fe(OH)3 --> FeCL3 + H2O FeCl3 --> Fe3+ + Cl- (inert), but Fe3+ = conj. acid of weak base and is VERY small but w. high charge --> high charge density --> all metals except for Na, K, Li, Ba can be weak acids but ONLY Fe(OH)3 AND Al(OH)3 can be salt hydrolysed bc. have very high charge densities: Fe3+ bc. charge density reacts w. wtr as complex ion bc. Fe bonds so well to H2O so will want to bond w. extra H2O: [Fe(H2O]6]3+ + H2O <--> [Fe(H2O)(OH)]2+ + H3O+ --> acidic sol. Happens bc. one H2O coordinately bonds to Fe 3+ bc. has such high charge density but O is most stable w. 2 bonds so lets go of one H so inst. forms OH- and the othr H reacts w. wtr into H3O+. Happens in othr metal compounds w.2+ as well, but to extremely small extent. Cations hydrolyse if is non-metal, if is metal dep.s on charge density Weak acid + strong base at 298K: H2CO3 + 2NaOH -->Na2CO3 + 2H2O Na2CO3 (aq) --> 2Na+ (=no hydrolysis) + CO3 2- (aq) CO3 2- + H2O <--> HCO3 - + OH- (aq) --> becomes basic sol. Weak acid + weak base at 298K: CH3COOH(aq) + NH3 --> NH4CH3COO --> NH4+ + CH3COO- = both strong conj.s => both hydrolyse wtr (NH4 + H2O <--> NH3 + H3O+; CH3COO- + H2O <--> CH3COOH + OH-); in THIS CASE resulting sol. is neutral as their pKa = pKb, but only neutral sol. IF H3O+ nd OH- are equal in amounts --> dep.s on how strong conj.s are in relation to each othr - if acid is slightly stronger will be slightly acidic nd vice versa; form salts where both conj.s carry out hydrolysis bc. both = rel.ly strong --> pH of sol. dep.s on rel. Ka nd Kb values of acids nd bases involved

Indicators

litmus dye --> acid blue to red, base red --> blue; best known acid-base indicator, derived from lichens, becomes pink in presence of acid, blue in alkalis, widely used to test for acids/alkalis but not as useful in distinguishing between strengths Indicators = chem. detectors, give info. abt. change in environment, most common in chem. = acid-base indicators changing colour reversibly acc. to H+ conc. in sol., happens bc. are weak acids nd bases whose conjugates have diff. colours - Δcolour = can be used to identify pH of substance, generally used as aq. sol.s or absorbed onto "test paper" Othr indicators (in Data Booklet) give diff. colours in diff. sol.s of acid/alkali (ex. methyl orange = colourless in acid, pink in alkali), many derived from nat.l substances ex. berry/flwr petal extract, common lab indicator = universal indicator (formed by mixing sev. indicators), changes colour many times across rnge of diff. acids/alkalis --> can be used to measuer H+ conc. on pH scale. Useful in acid-base titration in determining conc. of substance reacting w. standard sol. until equivalence p. reached where exactly neutralise each othr, convenient to have indicator distinctly changing colour when exactly neutralised Narrowe range indicators usually give more accurate readings than broad range, but always dep. on user's ability to interpret colours; pH meter/probe = often more accurate/objective than universal indicator paper, directly reads [H3O+] through spec. electrode, an record to accuracy of sev. decimal p.s, but must be calibrated before each use w. bufer sol. + standardised for temp. as pH is temp.-dep. Bc. [H3O+] x [OH-] gives constant value Kw these ions' conc.s must have inverse relation ships (in aq sol.s higher [H3O+] = lower [OH-], if we know either one we can calc. othr from Kw value - [H3O+] = Kw/[OH-] Rel. acidic strength of H halides incr. down gr. despite decr.ing molecular polarity - inst. xplaine by decr.ing bond strength as halogen atom incr.s in size; in othr cases ex. organic acids factors like electroneg., instability, inductive effects xplain patterns in acid strength Weak acid/base which has diff. coulour from its conj. so colour seen dep.s on if conj. has bonded or nore to H3O+ or OH- HIn(aq) (colour A) + H2O(l) <--> H3O+ + In-(aq) (colour B) If put in acidic sol. will be excess H3O+ to will be eqm shift left for more of colour A, when add base H3O+ is used up ny OH-s so eqm is shifted right to form more In- --> colour B --> Bc. weak acid in this case has Ka expression: Ka = [H3O+][In-]/[HIn]. At equivalence p. [HIn] = [In-] bc. eqm balence has been reached between acid nd its conj. base so cancel out --> Ka = [H3O+] --> pKa = pH --> indicators change colour when their pH is equal to pKa --> we can use pKa value of indicator to choose which indicators are suitable to which reactions. At this p. adding very small amount of acid/base --> shifts eqm as as above so causes pH to change colour = end-p., when pH is equal to pKa of indicator --> as diff. indicators have diff pKas change colours at diff. pH; indicator chosen should change oclour somewhere within pK range where jump occurs on pH curve. P. when pH changes colour = indicator's change point or end-point. Bc. indicators give visible change when pH changes can be used to identify eq. p.s in titrations bc. this is when pH changes most dramatically. Indicator = effective in signalling eq. of titration when its end-p. coincides w. pH at eq. p. - diff. indicators used for diff. titrations, dep.s on pH at eq. p. - 1. determine which combo of weak/strong acid/base will react, 2. deduce pH of salt sol. at eq. p. from nature of parent acid/base from salt hydrolysis, 3. choose indicator w. end.-p. in range of eq. p. via data tables Our eyes can identify distinct colours of one form of indicator when ratio of its conc. to conj. is abt. 10:1 , for colour to be observed from HIn to In- at end-p. ratio of these conc.s must change from 10:1 to 1:10 = range of 2 pH units => why is range of +/- 1 pH unit on either side of pKa value of indicator at which eye can definitely notice colour change occurring = end-p. range. When indicator used to detect eq. p in titration a few drops added from burette --> neutralisation reaction, exact vol. at which changes colour can be recorded as eq. p., diff. of just one drop of added sol. from burette should prod. dramatic change in colour at this p.

Brönsted-Lowry Theory of proton transfer (1923)

overcame Arrhenius's limitations; H atoms ionise into H+ + e-; acid = proton donor (whatever substance giving away H+), base = proton acceptor (whatever substance accepting H+); acid-base reactions must involve proton transfer, ex. HCl(aq) --> H+(aq) + Cl-(aq) --> H3o+(aq) + Cl-(aq) --> in order to be base must have lone pair of electrons (becomes obc. when draw Lewis structures); bc. is abt. H+ transfer is donation of it - donating can't occur in isolation - acid can only be acid if base present to accept H+ --> having pure HCl = nothing will happen, if HCl somes on skin skin will act as "base", proton acceptor, hence corrosive. Base must have reacted w. acid at first in order to be base. Ex. NH3(aq) + H2O(l) <--> NH4+(aq) + OH-(aq) H3O+ = hydro-/oxonium ion, makes sol. acidic (H+ singles DON'T EXIST in aq. sol.s but go to wtr molecules); ex. HCl + NH3 --> Cl- + NH4+ (wtr isn't shown but is ofc. present) OR HCl + NH4+ (bc. is in aq.so NH4 must exist as NH4+) + OH- --> H2O + NH4Cl bc. this = HCl + NH3 + HCl NH3(aq) + H2O(l) <--> NH4+(aq) + OH-(aq) = preferable way of writing, but othr way also OK