15.1 - Brønsted Acids and Bases

• The conjugate acid-base pair, which can be described as an acid and its conjugate base or a base and its conjugate acid, is an expansion of the Bronsted definition of acids and bases.

• There is a conjugate foundation of each Brønsted acid and a conjugate base of each Brønsted acid.

  • The chloride ion (Cl−), for example, is the conjugate base from HCl acid. The conjugated acid of bases H2O, is H3O+ (hydronium ion). + H2O + Cl− HCl + H2O

15.2 - The Acid-Base Properties of Water

• The ion-product constant, Kw, is defined as the product of the molar concentrations of H+ and OH ions at a given temperature.

• The hydrogen ion concentration is crucial for the study of acid-base reactions; its value shows the acidity or the fundamentality of the solution.

  • As very little water molecules are ionized, there remains virtually unchanged water concentration [H2O].

• The aquatic solution is called neutral whenever [H+] = [OH−]. Excess H+ ions and [H+]> [OH−] are present in an acidic solution.

• Excess hydroxide ions in a basic solution are [H+]<[OH−]. [H+] In practice, either H+ or OH− ions in solutions can be changed but neither can be independently varied.

  • The OH− concentration must be changed to [OH−] = Kw=[H+]= 1.0 − 10−14 1.0 | 10−6 = 1.0 − 10−8 M − to [OH−] Kw = 1.0 − H+ = 1.0 − 10−14 1.0 − 10−8 M.

15.3 - pH—A Measure of Acidity

• In 1909, Soren Sorensen created a more practical metric termed pH.

  • The negative logarithm of the hydrogen ion concentration is used to calculate the pH of a solution.

• There are reliable ways to estimate the solute concentration based on thermodynamics, but the details go beyond the scope of this text.

• The pH of a solution with a pH meter is measured in the laboratory

  • A number of common fluids are listed in Table 15.1 with pHs.

• Digestion is facilitated by a lower pH (high acidity) of gastric juices, while higher pH blood is required to transport oxygen.

15.4 - Strength of Acids and Bases

• Strong acids are strong electrolytes that are expected to ionize entirely in water for practical purposes.

  • The majority of acids are weak acids that only ionize to a limited extent in water.

  • Strong bases, like strong acids, are strong electrolytes that ionize in water.

  • Weak bases, like weak acids, are weak electrolytes. Ammonia is a weak base. It ionizes to a very limited extent in water

• Its conjugate base has no measurable strength if an acid is strong.

  • Therefore, the Cl−ion is an extremely faint base which is the conjugate base to strong HCl acid.

• H3O+ is the most potent and aqueous acid possible. Acids that are stronger than H3O+ react to H3O+ and its conjugate bases by water.

  • So, HCl, which is stronger than H3O+, fully reacts to H3O+ and Cl− with water: + H2O(l) + HCl(aq) + Cl− / HCl(aq) (aq)

• The OH−ion is the most powerful base in aqueous solution.

  • Bases that are stronger than OH− react to OH− and their conjugate acids with water.

  • Oxide ion (O2-), for instance, is a stronger base than OH- and therefore reacts completely like this with water: −(aq) + H2O(l) par 2OH− | O2−(aq) (aq)

15.5 - Weak Acids and Acid Ionization Constants

• Where Ka, the acid ionization constant, is the acid ionization equilibrium constant.

• The magnitude of Ka shows the acid's strength, as we've seen. The % ionization of acid is another indicator of its potency.

• Given the initial concentration and the Ka value, we can calculate the concentration of hydrogen ion or pH of the acid solution with equilibrium.

• Acid ionization constitutes an important category of aqueous solution chemical balance

  • we will develop a systemic procedure to solve such a problem, which will enable us also to understand the chemistry involved.

• Percent ionization which is defined as percent ionization = ionized acid concentration at equilibrium initial concentration of acid × 100% (15.11)

• The acid concentration under ionization for monoprotic acid HA is equal to the concentration of H+ ions or A− ions at equilibrium.

  • Thus we can write the% ionization in the form of [H+] [HA]0 = 100% ionization

15.6 - Weak Bases and Base Ionization Constants

• Because this reaction consumes extremely few water molecules in comparison to the total concentration of water, we can treat [H2O] as a constant.

  • As a result, we can write the base ionization constant (Kb), which is the ionization reaction's equilibrium constant.

• This substance is made up of Brønsted by the capacity of the lone pair to accept an H+ ion.

  • We follow the same procedure we used for weak acids to solve problems involving weak bases.

  • The main difference is that first, instead of [H+], we calculate [OH−].

15.7 - The Relationship Between the Ionization Constants of Acids and Their Conjugate Bases

• When two reactions are combined to form a third reaction, the third reaction's equilibrium constant is the product of the two added reactions' equilibrium constants.

• When two reactions add up, the balance constant for the third reaction is the equilibrium result of the balance constants for the two additional reactions.

  • So it always is true that KaKb = Kw (15.12) Expressing Equation (15.12) as Ka = Kw Kb Kb = Kw Ka For any combination of acid-base pairs.

15.8 - Diprotic and Polyprotic Acids

• Because diprotic and polyprotic acids can produce more than one hydrogen ion per molecule, their therapy is more complicated than monoprotic acids.

  • These acids ionize in a step-by-step fashion, losing one proton at a time.

• If Ka1 — bis Ka2 — for diprotic acids, it can be assumed that the H+ ion concentration is the product of an initial ionization phase only.

• In addition, the conjugate base concentration for ionization in the second stage is numerically equal to Ka2.

15.9 - Molecular Structure and the Strength of Acids

• The degree of ionization of the acid is influenced by two things.

  • One factor is the strength of the HX link—the stronger the connection, the more difficult it is for the HX molecule to break apart, and therefore the acid is weaker.

• The extent to which the acid is ionized influences two factors. One is an H-to-X bond strength — the stronger the bond.

• The harder it is for the HX molecule to break and, consequently, the weaker the acid.

• The polarity of the H umX bond is the other factor. The electron negative differences between H and X lead to a polar connection like

• The hypertension should be the strongest acid based on bond enthalpy as H + and I ions are most easily disrupted and formed. Secondly, look at the bond H-to-X polarity.

  • The polarity of the binder decreases in this series of acids from HF to HI as F is the most electronegative of halogens.

  • Because of the highest accumulation of positive and negative charges on H and F atoms HF is therefore the strongest acid based on bond polarity.

  • We therefore need to consider two competing factors to determine the strength of binary acids

15.10 - Acid-Base Properties of Salts

• Salts are powerful electrolytes that break down fully into ions when exposed to water.

• The reaction of a salt's anion or cation, or both, with water is known as salt hydrolysis.

• Salt that contains an alkaline metal ion or an alkaline earth metal ion (excluding Be2+)

  • Strong acid conjugated base are not generally hydrolysed and are presumably neutral.

• There are no significant acidic or basic properties to the hydrated Na+ ion.

  • But the acetate ion CH3COO− is the conjugate foundation of the CH3COOH faint acid and thus has an affinity with H+ ions.

15.11 - Acid-Base Properties of Oxides and Hydroxides

• Acidic oxides are transition metal oxides in which the metal has a high oxidation number.

  • Manganese(VII) oxide (Mn2O7) and chromium(VI) oxide (CrO3) are two well-known examples of metal oxides that react with water to form acids.

• If CO2 and H2O are affected, the reaction causes pure water to gradually reach a pH of approximately 5,5 when exposed to air (containing CO2).

  • Acid rain is largely responsible for the response between SO3 and H2O

  • The reactions between acidic oxides and the basic oxides and acids are similar to normal acid-based reactions as salt and water are the products: CO2(g) 2NaOH(aq) Na2CO3(aq) H2O(l) salt water base acidic oxide. H2O(l) basic water oxide salt oxide Ba(NO3)2(aq) H2O(l)

15.12 - Lewis Acids and Bases

• G. N. Lewis, an American scientist, proposed such a definition in 1932. He described a Lewis base as a material capable of donating a pair of electrons.

  • A Lewis acid is an atom or molecule that can receive two electrons.

• A Lewis acid-base reaction involves donation from one species to another of a pair of electrodes.

  • A salt and water reaction like that is not producing.

• It's more general than other definitions that the Lewis concept has significance.

• Lewis acid-base reactions include many non-Brønsted acid reactions.

  • For instance, consider the reaction to an adduct compound between boron trifluoride (BF3) and ammonia