Weak acids defined

The first case simply involves the ionization of a weak acid and is governed by the expression that defines Ka for acetic acid ... 

In order to obtain good yields from a Weiss reaction sequence, the H+-concentration has to be adjusted properly in the reaction mixture. The reaction is usually carried out in a buffered, weakly acidic or weakly basic solution. By the Weiss reaction simple starting materials are converted into a complex product of defined stereochemistry. There is no simpler procedure for the synthesis of the l,5-c -disubstituted bicyclo[3.3.0]octane skeleton it has for example found application in the synthesis of polyquinanes. ... 

A large number of other reactions are also possible, e.g., hydration of the various substances, or (if the solute is a salt of a weak acid or base) hydrolysis, but in all cases the concentration of each molecular species is defined by the total amount of solute in a given mass of solution, and the ionisation proceeds as if all the other reactions did not occur at all (cf. 158). 

As Figure 17-6 indicates, only a fraction of weak acid molecules undergoes proton transfer to form hydronium ions. Equation defines the percent ionization of any acid solution to be the percentage of the initial concentration... 

To protect a solution against pH variations, a major species in the solution must react with added hydronium ions, and another major species must react with added hydroxide ions. The conjugate base of a weak acid will react readily with hydronium ions, and the weak acid itself will react readily with hydroxide ions. This means that a buffer solution can be defined in terms of its composition. 

Arrhenius postulated in 1887 that an appreciable fraction of electrolyte in water dissociates to free ions, which are responsible for the electrical conductance of its aqueous solution. Later Kohlrausch plotted the equivalent conductivities of an electrolyte at a constant temperature against the square root of its concentration he found a slow linear increase of A with increasing dilution for so-called strong electrolytes (salts), but a tangential increase for weak electrolytes (weak acids and bases). Hence the equivalent conductivity of an electrolyte reaches a limiting value at infinite dilution, defined as... 

When placed in water, some acids and bases completely ionize, or dissociate into their ions—but others do not. Acids and bases that completely ionize are called strong acids and strong bases. Strong acids are defined as acids that have a pH of 0-4. Strong bases have pH values of 10-14. On the other hand, weak acids and weak bases do not disassociate completely in water. This leads to a pH value that is closer to neutral, because some of the hydrogen ions are still attached to other atoms, decreasing the hydrogen ion concentration. 

It is convenient to summarize the various reactions in a box diagram, such as Fig. 4.1 [17,275,280], illustrated with the equilibria of the weak base, propranolol. In Fig. 4.1 is an equation labeled pA °et. This constant refers to the octanol pKa, a term first used by Scherrer [280]. When the concentrations of the uncharged and the charged species in octanol are equal, the aqueous pH at that point defines p which is indicated for a weak acid as... 

Electrolytes are defined as substances whose aqueous solutions conduct electricity due to the presence of ions in solution. Acids, soluble bases and soluble salts are electrolytes. Measuring the extent to which a substance s aqueous solution conducts electricity is how chemists determine whether it is a strong or weak electrolyte. If the solution conducts electricity well, the solute is a strong electrolyte, like the strong acid, HC1 if it conducts electricity poorly, the solute is a weak electrolyte, like the weak acid, HF. 

But we need to be careful when talking about the magnitudes of Consider the case of sodium ethanoate dissolved in dilute mineral acid the reaction occurring is, in fact, the reverse of that in Equation (4.45), with a proton and carboxylate anion associating to form undissociated acid. In this case, = 1 mol before the reaction occurs, and its value decreases as the reaction proceeds. In other words, we need to define our reaction before we can speak knowledgeably about it. We can now rewrite our question, asking Why is < 1 for a weak acid ... 

Equilibria involving acids and bases are discussed from within the Lowry-Br0nsted theory, which defines an acid as a proton donor and a base as a proton acceptor (or abstracter ). The additional concept of pH is then introduced. Strong and weak acids are discussed in terms of the acidity constant Ka, and then conjugate acids and bases are identified. 

A buffer comprises (1) a weak acid and a salt of that acid, (2) a weak base and a salt of that base, or (3) it may contain an acid salt. We define an acid-base buffer as a solution whose pH does not change after adding (small amounts of) a strong acid or base . Sodium ascorbate is a favourite buffer in the food industry. 

In the process of a weak acid or weak base neutralization titration, a mixture of a conjugate acid-base pair exists in the reaction flask in the time period of the experiment leading up to the inflection point. For example, during the titration of acetic acid with sodium hydroxide, a mixture of acetic acid and acetate ion exists in the reaction flask prior to the inflection point. In that portion of the titration curve, the pH of the solution does not change appreciably, even upon the addition of more sodium hydroxide. Thus this solution is a buffer solution, as we defined it at the beginning of this section. 

The first equation defines the ionization constant of water at 25 °C (we omit the sign of the charges to simplify notation). The second is the same as Eq. (2.6.3), while the third is the conservation of the total initial concentration of the (weak) acid Nj- (we assume that there is no change in volume during the titration, hence this is the same as the conservation of the total number of acid molecules). The fourth equation is the electroneutrality condition, where [iV ] is the concentration of the added (strong) base. 

For a monobasic weak acid, HA, the thermodymanic dissociation constant (fiT ) is defined as follows ... 

The answer is D. Weak acids like lactic acid never completely dissociate in solution and are thus defined by the property that at least some of the protonated (undissociated acid) form and the unprotonated (conjugate base) form of the acid are present at all concentrations and pH conditions. The indicated of 5.2 is consistent with the idea that the lactate anion retains a strong affinity for protons, a hallmark of a weak acid. The lactate anion is highly water-soluble. All weak acids obey the Henderson-Hasselbalch equation. 

J. Kendall, J. E. Booze, and J. C. Andrews 3 have shown that the formation of hydrates, in the sense of water of crystallization, with the weak acids very seldom occurs, and when hydrates are formed, the acid has the amphoteric character of a phenol. There is also a regular increase in the tendency of an acid to form hydrates, as the strength of the acid increases, until, with the strong acids, well-defined stable hydrates appear. The complexity and stability of the hydrates increase with the strength of the acid. These facts are in harmony with the weak acid nature of water. 

Being the opposite of an acid, a base will be defined as a compound that has a tendency to combine with protons. In this definition the base9 in an alkaline solution is the OH ion. This ion is one of the strongest bases known to exist. When combining with a proton it forms water, that itself is a weak base, since it is able to add one more proton to form an OHJ, hydronium ion, and a weak acid at the same time, since water can dissociate into OH and H+ ions. Water being a base, too, the actual dissociation reaction will be... 

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