# Question cc730

Mar 3, 2017

I am not entirely sure what you asking here. If this inadequate, say the word, and someone will give it another go..........

#### Explanation:

In aqueous solution, i.e. in water, we conceive that the solvent water undergoes the following equilibrium:

$2 {H}_{2} O \left(l\right) r i g h t \le f t h a r p \infty n s {H}_{3} {O}^{+} + H {O}^{-}$

We can very precisely measure the extent of this equilibrium by conductivity experiments, and we find that, under standard conditions, that this acid-base reaction, this autoprotolysis may be quantified according to:

${K}_{w} = \left[{H}_{3} {O}^{+}\right] \left[H {O}^{-}\right] = {10}^{- 14}$

Any species added to solution that raises $\left[{H}_{3} {O}^{+}\right]$ (and this is conceived of as the $\text{hydronium ion}$, the characteristic cation of the water solvent) is an $\text{acid}$, and a species that raises [""^(-)OH] is a base (i.e. the characteristic anion). At equilibrium, however, the given relationship operates.

The strength of an acid is determined by the extent its equilibrium lies to the right in the given equation:

$H A \left(a q\right) + {H}_{2} O \left(l\right) r i g h t \le f t h a r p \infty n s {H}_{3} {O}^{+} + {A}^{-}$.

For strong acids, such as HX, (X=Cl, Br, I;X!=F), ${H}_{2} S {O}_{4}$, $H C l {O}_{4}$, this equilibrium lies almost entirely to the right. And thus $1 \cdot m o l \cdot {L}^{-} 1$ solutions of $H C l , H B r$, are $1 \cdot m o l \cdot {L}^{-} 1$ with respect to $\left[{H}_{3} {O}^{+}\right]$, AND to $\left[{X}^{-}\right]$, which is the so-called conjugate base of the original acid. Note that the definition specifies the interaction of the acid with the solvent. You can buy gas cylinders of $H C l$; in water it participates in the given equilibrium.

For weaker acids, such as $H F$, in aqueous solution some concentration of $\left[H F\right]$ remains, because the protolysis reaction does not proceed so far. Are you following me?

Mar 3, 2017

Here's what I get.

#### Explanation:

Weak acid dissociation

A weak acid is an acid that does not dissociate completely in water.

Most of the molecules remain undissociated.

Consider an aqueous solution of a weak acid $\text{HA}$.

$\text{HA + H"_2"O" ⇌ "H"_3"O"^"+" + "A"^"-}$

Every time a molecule of $\text{HA}$ dissociates, it forms 1 $\text{H"_3"O"^"+}$ ion and 1 $\text{A"^"-}$ ion.

Thus,

["H"_3"O"^"+"] = ["A"]^"-"#

Strong acid dissociation

A strong acid is an acid that dissociates completely in water.

Consider an aqueous solution of a strong acid $\text{HX}$.

$\text{HX + H"_2"O" → "H"_3"O"^"+" + "X"^"-}$

Every time a molecule of $\text{HX}$ dissociates, it forms 1 $\text{H"_3"O"^"+}$ ion and 1 $\text{X"^"-}$ ion.

However, almost every molecule of $\text{HX}$ dissociates.

If we have 1 mol of $\text{HX}$, we will get 1 mol of $\text{H"_3"O"^"+}$ ions and 1 mol of $\text{X"^"-}$ ions.

Thus, the concentration of $\text{H"_3"O"^"+}$ ions is the same as the initial concentration of the $\text{HX}$.

or

${\left[\text{H"_3"O"^"+"] = ["HX}\right]}_{0}$