# PH

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pH = -log<sub>10</sub> [H<sup>+</sup>] | pH = -log<sub>10</sub> [H<sup>+</sup>] | ||

− | The pH scale is a measure of | + | The pH scale is a measure of [[Hydrogen_ion|hydrogen ion]] concentration that eliminates dealing with large powers of 10 and compresses a large range of concentrations onto a more convenient scale, between 1 and 14 as show in the figure below: |

[[Image:PH.png]]<br> | [[Image:PH.png]]<br> | ||

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=== Acid Dissociation<br> === | === Acid Dissociation<br> === | ||

− | Strong [[Acid|acids]] are considered to be completely dissociated into [[Ion|ions]] in dilute solutions. However, weak [[Acid|acids]] (or [[Base|bases]]) are | + | Strong [[Acid|acids]] are considered to be completely dissociated into [[Ion|ions]] in dilute solutions. However, weak [[Acid|acids]] (or [[Base|bases]]) are partially dissociated in solution, and thus an [[Equilibrium|equilibrium]] is established between the [[Ion|ions]] and the undissociated [[Molecules|molecules]].<br> |

This equilibrium can be represented by the equation: <br> | This equilibrium can be represented by the equation: <br> | ||

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[HA]<=>[H<sup>+</sup>]+[A<sup>-</sup>]. <br> | [HA]<=>[H<sup>+</sup>]+[A<sup>-</sup>]. <br> | ||

− | Where: [HA] is the concentration of undissociated molecules; [H+] is the concentration of hydrogen ions; and [A-] is the concentration of the conjugate base.<br> | + | Where: [HA] is the [[Concentration|concentration]] of undissociated molecules; [H+] is the concentration of hydrogen ions; and [A-] is the concentration of the conjugate [[Base|base]].<br> |

− | This equation can be rearranged to give a K<sub>a</sub> value, which is a measure of how strong an acid is. | + | This equation can be rearranged to give a K<sub>a</sub> value, which is a measure of how strong an [[Acid|acid]] is. |

K<sub>a</sub>= [H<sup>+</sup>][A<sup>-</sup>]/[HA] | K<sub>a</sub>= [H<sup>+</sup>][A<sup>-</sup>]/[HA] | ||

− | Stronger acids will dissociate more and will have a higher K<sub>a</sub> value | + | Stronger acids will dissociate more and will have a higher K<sub>a</sub> value <ref>Elliot WH &amp;amp;amp;amp;amp; Elliot DC, 2009, Biochemistry and Molecular Biology p.38; Oxford</ref>. |

− | Use of K<sub>a</sub> values is not very useful, as the differences in dissociation are massive between strong and weak acids, and the values calculated can vary by many orders of magnitude. As is done with the pH scale, we take the negative log of the K<sub>a</sub> value to give the pK<sub>a</sub> value | + | Use of K<sub>a</sub> values is not very useful, as the differences in dissociation are massive between strong and weak acids, and the values calculated can vary by many orders of magnitude. As is done with the pH scale, we take the negative log of the K<sub>a</sub> value to give the pK<sub>a</sub> value <ref>Elliot WH &amp;amp;amp;amp;amp; Elliot DC, 2009, Biochemistry and Molecular Biology p.39; Oxford</ref>. |

− | pK<sub>a</sub>=-log<sub>10</sub>K<sub>a</sub> | + | pK<sub>a</sub>=-log<sub>10</sub>K<sub>a</sub> |

− | When: pK<sub>a</sub>=pH; the compound is said to be at it's equivalence point, as the concentration of H+ is equal to the concentration of HA. | + | When: pK<sub>a</sub>=pH; the [[Compound|compound]] is said to be at it's equivalence point, as the [[Concentration|concentration]] of H+ is equal to the concentration of HA. |

=== References === | === References === | ||

<references /> | <references /> |

## Latest revision as of 14:42, 24 November 2013

The negative logarithm of the hydrogen ion concentration, the pH, is expressed as follows:

pH = -log_{10} [H^{+}]

The pH scale is a measure of hydrogen ion concentration that eliminates dealing with large powers of 10 and compresses a large range of concentrations onto a more convenient scale, between 1 and 14 as show in the figure below:

### Acid Dissociation

Strong acids are considered to be completely dissociated into ions in dilute solutions. However, weak acids (or bases) are partially dissociated in solution, and thus an equilibrium is established between the ions and the undissociated molecules.

This equilibrium can be represented by the equation:

[HA]<=>[H^{+}]+[A^{-}].

Where: [HA] is the concentration of undissociated molecules; [H+] is the concentration of hydrogen ions; and [A-] is the concentration of the conjugate base.

This equation can be rearranged to give a K_{a} value, which is a measure of how strong an acid is.

K_{a}= [H^{+}][A^{-}]/[HA]

Stronger acids will dissociate more and will have a higher K_{a} value ^{[1]}.

Use of K_{a} values is not very useful, as the differences in dissociation are massive between strong and weak acids, and the values calculated can vary by many orders of magnitude. As is done with the pH scale, we take the negative log of the K_{a} value to give the pK_{a} value ^{[2]}.

pK_{a}=-log_{10}K_{a}

When: pK_{a}=pH; the compound is said to be at it's equivalence point, as the concentration of H+ is equal to the concentration of HA.