> See also: > - [[Acids and Bases]] # Relative Strengths of Acids & Bases > When determining the acidity of a molecule, simply look at the stability of the conjugate base Its important to compare the relative strengths of acids rather than using absolute hierarchies(?) > - *Strong Acids* completely dissolves/dissociates in water > - *Weak Acids* do not dissolve nearly as much as strong acids | Classification | Range | | --- | --- | | Very Strong Acids | $pK_a <1$ | | Moderately Strong Acids | $pK_a = 1-3$ | | Weak Acids | $pK_a = 3-5$ | | Very Weak Acids | $pK_a = 5-15$ | | Extremely Weak Acids | $pK_a > 15$ | ![[Pasted image 20220831092254.png|400]] ![[Pasted image 20220831092300.png|400]] ## Measurements of [[Acids and Bases|Acidity and Basicity]] (pH & pOH) | Classification | Relative Ion Concentrations | pH at $25 \degree C$ | |:---:|:---:|:---:| | Acidic | $[H_3O^+] > [OH^-]$ | pH < 7 | | Neutral | $[H_3O^+] = [OH^-]$ | pH = 7 | | Basic | $[H_3O^+] < [OH^-]$ | pH > 7 | > [!info] **Calculating pH & pOH** > The pH of a solution is calculated using the concentration of *hydronium ions*: > $pH = -\log [H_3O^+]$ > --- > The pH of a solution is calculated using the concentration of *hydroxide ions*: > $pOH = -\log [OH^-]$ These functions are essentially the same, they are both negative [[logarithms]] using a specified concentration. The generalized version of this relationship is known as a p-function: $pX = -\log[X]$ Using the properties of [[logarithms]], this equation can be reversed to form the following equation: $[X] = 10^{-pX}$ ### Measurement Scales > [!info] > Both of these scales **are NOT fixed**, meaning that pH and pOH values can technically be below 0 or above 14. > [!example]- pH Chart With Common Examples > ![[Pasted image 20220716195206.png|600]] - [[Buffer Solutions]] [[Factors that Stabilize Negative Charges]]