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Biomedical Science Health and Society, Malmö University Sergey Shleev Chemistry “Acids, bases, pH, pK, pI”

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Antoine Lavoisier (1743 - 1794) Pioneer of analytical chemistry and chemical nomenclature ”One of the first defenition of acids” 1780 Nitrogen (N 2 ) – kväve – (Greek: “no life”) Hydrogen (H 2 ) – väte – gas which is “water-former“ (Greek) Oxygen (O 2 ) – syrgas – gas which is "acid-former“ (Greek: “becoming sharp”) H 2 SO 4 HClO 4

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Simple Complex Nonmetals [S, P] Metals [K, Na, Ca] Oxides acidic [SO 2, SO 3, P 2 O 5, ] basic [K 2 O, Na 2 O, CaO] Acids Bases [KOH, NaOH, Ca(OH) 2 ] Salts [KCl,, NaSO 4 ] oxoacids [HPO 3, H 4 P 2 O 7, H 3 PO 4 H 2 SO 3, H 2 SO 4 ] nonoxo [HCl, HF, HCN, H 2 S] Hydride Metals [KH, NaH, CaH 2 ] Nonmetals [PH 3, SiH 4 ] Inorganic substances

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Svante August Arrhenius (1859 - 1927) ”Theory of electrolytic dissociation” – 1887 Nobel Price in Chemistry – 1903 Arrhenius theory that describes aqueous solutions in terms of acids (which dissociate to give hydrogen ions) and bases (which dissociate to give hydroxyl ions); the product of an acid and a base is a salt and water This theory is still involved in our modern understanding of electrolytes, electrical conductivity of solutions, etc

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HCl NH 4 OH Whit e smok e HCl + NH 3 = NH 4 Cl ? HCl + NH 4 OH = NH 4 Cl + H 2 O ?? Al(OH) 3 or H 3 AlO 3 ??? [Al(H 2 O) 6 ] 3+ electrical conductivity of melts?

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Johannes Nicolaus Br ø nsted (1879–1947) Thomas Martin Lowry (1874–1936) 1923- formulation of the protonic definition of acids and bases, both in solution and in gas phase

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Gilbert Newton Lewis (1875–1946) 1923 - the electron-pair theory of acid-base reactions the donation of electron pairs from bases and the acceptance by acids, rather than protons or other bonded substances The theory spans both aqueous and non-aqueous reactions

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Inorganic substances Simple Complex nonmetals [S, P] metals [K, Na, Ca] Oxides acidic [SO 2, SO 3, P 2 O 5, ] basic [K 2 O, Na 2 O, CaO] Basic hydroxides [KOH, NaOH] Salts [NaSO 4 ] Acidic hydroxides [HPO 3, H 4 P 2 O 7, H 3 PO 4 H 2 SO 3, H 2 SO 4 ] Hydride metals [KH, NaH] Amphoteric [Al] Amphoteric [Al 2 O 3 ] Amphoteric hydroxides [Al(OH) 3 ] Nonmetals [PH 3, SiH 4 ] Halides Metalic [KCl, NaF] Nonmetalic [HCl, HF] Chelates [K 3 Fe(CN) 6 ]

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Arrhenius Br ø nsted -Lowry

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Strong acid Week acid

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Dissociation of an acid can be written in several possible ways: HA + A + H + HA A - + H + HA - A 2- + H + Note that in some cases the conjugate base (A, A -, or A 2- ) has a negative charge and in other cases it does not, but in all cases it has one less positive charge than the acid. For convenience, we will always write such reactions as HA H + + A - The equilibrium constant (Ka) for the dissociation of a weak acid (often called the dissociation constant):

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Mathematics “Logarithms, exponents, and quadratic equations”

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exponents observed variable constant controlled variable Euler's number

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Arrhenius law (equation) activation energy (T independent constant; (kJ mol -1 ) frequency or pre-exponential factor temperatue (K) gas constant (8.3155 J K -1 mol -1 ) rate constant of chemical reaction

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Logarithms Natural logarithm ln(x) – the inverse function to exponential (e x ) log(x) – the inverse function to 10 x So…

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pK a «acid dissociation constant» pK b «base dissociation constant»

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“self-ionization constant of water” hydronium hydroxide

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pH «power of hydrogen» pOH

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quadratic equations

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Henderson–Hasselbalch equation

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Buffer capacity

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pI «isoelectic point» the pH at which a particular molecule or surface carries no net electrical charge For glycine, an amino acid with only one amine and one carboxyl group 6.062.359.78

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Asp?1.999.90 Lys?2.169.06 pI pK 1 pK 2 pK s 3.90 10.54 (9.6) (2.8)

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The relative concentrations of the three forms of glycine as a function of pH

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