# percent ionization calculator

Take 1.0 L of solution to have the quantities on a mole basis. Water also exerts a leveling effect on the strengths of strong bases. Substituting the equilibrium concentrations into the equilibrium expression and making the assumption that (0.120 − x) ≈ 0.120 gives: $\frac{\left[\text{Fe}{\left({\text{H}}_{2}\text{O}\right)}_{5}{\left(\text{OH}\right)}^{+}\right]\left[{\text{H}}_{3}{\text{O}}^{+}\right]}{\left[\text{Fe}{\left({\text{H}}_{2}\text{O}\right)}_{6}^{2+}\right]}=\frac{\left(x\right)\left(x\right)}{\left(0.120-x\right)}\approx \frac{\left(x\right)\left(x\right)}{0.120}=1.6\times {10}^{-7}$. A solution consisting of a certain concentration of the powerful acid HCl, hydrochloric acid, will be "more acidic" than a solution containing a similar concentration of acetic acid, or plain vinegar. Strong bases react with water to quantitatively form hydroxide ions. The initial concentration of ${\text{H}}_{3}{\text{O}}^{+}$ is its concentration in pure water, which is so much less than the final concentration that we approximate it as zero (~0). Substituting the equilibrium concentrations into the equilibrium expression and making the assumption that (0.0092 − x) ≈ 0.0092 gives: $\frac{\left[{\text{H}}_{3}{\text{O}}^{+}\right]\left[{\text{ClO}}^{-}\right]}{\left[\text{HClO}\right]}=\frac{\left(x\right)\left(x\right)}{\left(0.0092-x\right)}\approx \frac{\left(x\right)\left(x\right)}{0.0092}=3.5\times {10}^{-8}$. Explain why the neutralization reaction of a strong acid and a weak base gives a weakly acidic solution. Identify the strong Brønsted-Lowry acids and strong Brønsted-Lowry bases. The acidity increases as the electronegativity of the central atom increases. Thus, the order of increasing acidity (for removal of one proton) across the second row is CH4 < NH3 < H2O < HF; across the third row, it is SiH4 < PH3 < H2S < HCl (see Figure 7). Finally, use the other equilibrium to find the other concentrations. 2. This reaction also forms OH−, which causes the solution to be basic. The extent to which a base forms hydroxide ion in aqueous solution depends on the strength of the base relative to that of the hydroxide ion, as shown in the last column in Figure 2. Compounds that are weaker acids than water (those found below water in the column of acids) in Figure 3 exhibit no observable acidic behavior when dissolved in water. $\left[\text{H}_{2}\text{O}\right]\gt\left[\text{C}_{6}\text{H}_{4}\text{OH}\left(\text{CO}_{2}\text{H}\right)\right]\gt\left[\text{H}^{+}\right]\text{O}\gt\left[\text{C}_{6}\text{H}_{4}\text{OH}\left(\text{CO}_{2}\right)^{-}\right]\gt\left[\text{C}_{6}\text{H}_{4}\text{O}\left(\text{CO}_{2}\text{H}\right)^{-}\right]\gt\left[\text{OH}^{-}\right]$, 2. Similarly we can obtain the formula for a weak base like B whose conjugate acid is BH+. For the reaction of an acid HA: $\text{HA}\left(aq\right)+{\text{H}}_{2}\text{O}\left(l\right)\rightleftharpoons {\text{H}}_{3}{\text{O}}^{+}\left(aq\right)+{\text{A}}^{-}\left(aq\right)$. In a solution containing a mixture of NaH2PO4 and Na2HPO4 at equilibrium, [OH−] = 1.3 × 10−6M; $\left[{\text{H}}_{2}{\text{PO}}_{4}^{-}\right]=0.042M$; and $\left[{\text{HPO}}_{4}^{2-}\right]=0.341M$. Because the ratio includes the initial concentration, the percent ionization for a solution of a given weak acid varies depending on the original concentration of the acid, and actually decreases with increasing acid concentration. Because acids are proton donors, in everyday terms, you can say that a solution containing a "strong acid" (that is, an acid with a high propensity to donate its protons) is "more acidic."