unit:PH

Type

Description

In chemistry the unit $pH$ , also referred to as $acidity$ or $basicity$ , is the negative logarithm (base 10) of the concentration of free protons (or hydronium ions). The definition of $p H$ in terms of hydrogen ions in solution is: $pH = - \log_{10} (a_{H^{+}}) \equiv - \log_{10} ([H^{+}])$ Where $a_{H^{+}}$ is the equilibrium molar concentration of $H^{+}$ in the solution, the activity of the hydrogen ion in the solution. This definition is appropriate for concentrations equal to, or less than $1 m o l / l$ , where $a H + \equiv [H +]$ , that is, $1 m o l / L H C l$ has a $p H$ of zero. To relate this to standard molality ( $b^{\circ}$ ), typically taken as $1 m o l / k g$ , consideration is given to the activity ( $a_{H^{+}}$ ) of the hydrogen ions. The activity can be expressed as: $a_{H^{+}} = γ_{H^{+}} \times m_{H^{+}}$ Where, $γ_{H^{+}}$ is the activity coefficient, which adjusts the molality to account for non-ideal behavior due to interactions between ions in the solution. $m_{H^{+}}$ is the molality of hydrogen ions in the solution relative to the standard molality, expressed in $m o l / k g$ . The expansion of $p H$ then becomes: $pH = - l o g_{10} (m_{H +} \times γ_{H^{+}})$ This definition is relevant in more concentrated solutions or when precise thermodynamic calculations are required. It reflects how the properties of the solution deviate from ideal behavior and provides a more accurate understanding of the $p H$ under various conditions. While $p H$ is a universally recognized scale for expressing hydrogen ion activity, its appropriateness and accuracy can diminish under conditions of extremely high ionic strength, non-aqueous environments, high temperatures, or very high or low $p H$ values. In such cases, alternative measurement strategies might be required to obtain meaningful and accurate descriptions of acidity or basicity.

Properties

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Annotations

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dcterms:description

In chemistry the unit

pH

, also referred to as

acidity

basicity

, is the negative logarithm (base 10) of the concentration of free protons (or hydronium ions). The definition of

p H

in terms of hydrogen ions in solution is:

pH = - \log_{10} (a_{H^{+}}) \equiv - \log_{10} ([H^{+}])

Where

a_{H^{+}}

is the equilibrium molar concentration of

H^{+}

in the solution, the activity of the hydrogen ion in the solution.

This definition is appropriate for concentrations equal to, or less than

1 m o l / l

, where

a H + \equiv [H +]

, that is,

1 m o l / L H C l

has a

p H

of zero.

To relate this to standard molality (

b^{\circ}

), typically taken as

1 m o l / k g

, consideration is given to the activity (

a_{H^{+}}

) of the hydrogen ions.

The activity can be expressed as:

a_{H^{+}} = γ_{H^{+}} \times m_{H^{+}}

Where,

γ_{H^{+}}

is the activity coefficient, which adjusts the molality to account for non-ideal behavior due to interactions between ions in the solution.

m_{H^{+}}

is the molality of hydrogen ions in the solution relative to the standard molality, expressed in

m o l / k g

The expansion of

p H

then becomes:

pH = - l o g_{10} (m_{H +} \times γ_{H^{+}})

This definition is relevant in more concentrated solutions or when precise thermodynamic calculations are required. It reflects how the properties of the solution deviate from ideal behavior and provides a more accurate understanding of the

p H

under various conditions.

While

p H

is a universally recognized scale for expressing hydrogen ion activity, its appropriateness and accuracy can diminish under conditions of extremely high ionic strength, non-aqueous environments, high temperatures, or very high or low

p H

values. In such cases, alternative measurement strategies might be required to obtain meaningful and accurate descriptions of acidity or basicity.

qudt:informativeReference

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View as: CSV

Work in progress

RDF/XML

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  In chemistry the unit $\textit{pH}$, also referred to as $\textit{acidity}$ or $\textit{basicity}$,
   is the negative logarithm (base 10) of the concentration of free protons (or hydronium ions).
  The definition of $pH$ in terms of hydrogen ions in solution is:

  $$\text{pH}=-\log_{10}(a_{H^+})\equiv-\log_{10}\left(\left[H^+\right]\right)$$

  Where $a_{H^+}$ is the equilibrium molar concentration of $H^+$ in the solution, the activity of
   the hydrogen ion in the solution.
  $$$$
  This definition is appropriate for concentrations equal to, or less than $1\ mol/l$, 
   where $aH+ \equiv [H+]$, that is, $1\ mol/L\ HCl$ has a $pH$ of zero.
  $$$$
  To relate this to standard molality ($b^\circ$), typically taken as $1 \ mol/kg$, 
   consideration is given to the activity ($a_{H^+}$) of the hydrogen ions.
  $$$$
  The activity can be expressed as:

  $$a_{H^+} = \gamma_{H^+} \times m_{H^+}$$ 

  Where, $\gamma_{H^+}$ is the activity coefficient, which adjusts the molality to account for
   non-ideal behavior due to interactions between ions in the solution.
  $m_{H^+}$ is the molality of hydrogen ions in the solution relative to the standard molality, 
   expressed in $mol/kg$.
  $$$$
  The expansion of $pH$ then becomes:

  $$\text{pH} = -log_{10}\left(m_{H+}\times\gamma_{H^+}\right)$$

  $$$$
  This definition is relevant in more concentrated solutions or when precise thermodynamic calculations are required. 
  It reflects how the properties of the solution deviate from ideal behavior and provides a more accurate understanding of the $pH$ under various conditions.
  $$$$
  While $pH$ is a universally recognized scale for expressing hydrogen ion activity,
  its appropriateness and accuracy can diminish under conditions of extremely high
  ionic strength, non-aqueous environments, high temperatures, or very high or low $pH$ values.
  In such cases, alternative measurement strategies might be required to obtain meaningful and accurate descriptions of acidity or basicity.
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TURTLE

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  rdf:type <http://qudt.org/schema/qudt/Unit> ;
  <http://purl.org/dc/terms/description> """
  In chemistry the unit $\\textit{pH}$, also referred to as $\\textit{acidity}$ or $\\textit{basicity}$,
   is the negative logarithm (base 10) of the concentration of free protons (or hydronium ions).
  The definition of $pH$ in terms of hydrogen ions in solution is:

  $$\\text{pH}=-\\log_{10}(a_{H^+})\\equiv-\\log_{10}\\left(\\left[H^+\\right]\\right)$$

  Where $a_{H^+}$ is the equilibrium molar concentration of $H^+$ in the solution, the activity of
   the hydrogen ion in the solution.
  $$$$
  This definition is appropriate for concentrations equal to, or less than $1\\ mol/l$, 
   where $aH+ \\equiv [H+]$, that is, $1\\ mol/L\\ HCl$ has a $pH$ of zero.
  $$$$
  To relate this to standard molality ($b^\\circ$), typically taken as $1 \\ mol/kg$, 
   consideration is given to the activity ($a_{H^+}$) of the hydrogen ions.
  $$$$
  The activity can be expressed as:

  $$a_{H^+} = \\gamma_{H^+} \\times m_{H^+}$$ 

  Where, $\\gamma_{H^+}$ is the activity coefficient, which adjusts the molality to account for
   non-ideal behavior due to interactions between ions in the solution.
  $m_{H^+}$ is the molality of hydrogen ions in the solution relative to the standard molality, 
   expressed in $mol/kg$.
  $$$$
  The expansion of $pH$ then becomes:

  $$\\text{pH} = -log_{10}\\left(m_{H+}\\times\\gamma_{H^+}\\right)$$

  $$$$
  This definition is relevant in more concentrated solutions or when precise thermodynamic calculations are required. 
  It reflects how the properties of the solution deviate from ideal behavior and provides a more accurate understanding of the $pH$ under various conditions.
  $$$$
  While $pH$ is a universally recognized scale for expressing hydrogen ion activity,
  its appropriateness and accuracy can diminish under conditions of extremely high
  ionic strength, non-aqueous environments, high temperatures, or very high or low $pH$ values.
  In such cases, alternative measurement strategies might be required to obtain meaningful and accurate descriptions of acidity or basicity.
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JSON

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    ,"applicable system":"sou:PLANCK" 
    ,"applicable system":"sou:SI" 
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    ,"description":"\n  In chemistry the unit $\\textit{pH}$, also referred to as $\\textit{acidity}$ or $\\textit{basicity}$,\n   is the negative logarithm (base 10) of the concentration of free protons (or hydronium ions).\n  The definition of $pH$ in terms of hydrogen ions in solution is:\n\n  $$\\text{pH}=-\\log_{10}(a_{H^+})\\equiv-\\log_{10}\\left(\\left[H^+\\right]\\right)$$\n\n  Where $a_{H^+}$ is the equilibrium molar concentration of $H^+$ in the solution, the activity of\n   the hydrogen ion in the solution.\n  $$$$\n  This definition is appropriate for concentrations equal to, or less than $1\\ mol\/l$, \n   where $aH+ \\equiv [H+]$, that is, $1\\ mol\/L\\ HCl$ has a $pH$ of zero.\n  $$$$\n  To relate this to standard molality ($b^\\circ$), typically taken as $1 \\ mol\/kg$, \n   consideration is given to the activity ($a_{H^+}$) of the hydrogen ions.\n  $$$$\n  The activity can be expressed as:\n\n  $$a_{H^+} = \\gamma_{H^+} \\times m_{H^+}$$ \n\n  Where, $\\gamma_{H^+}$ is the activity coefficient, which adjusts the molality to account for\n   non-ideal behavior due to interactions between ions in the solution.\n  $m_{H^+}$ is the molality of hydrogen ions in the solution relative to the standard molality, \n   expressed in $mol\/kg$.\n  $$$$\n  The expansion of $pH$ then becomes:\n\n  $$\\text{pH} = -log_{10}\\left(m_{H+}\\times\\gamma_{H^+}\\right)$$\n\n  $$$$\n  This definition is relevant in more concentrated solutions or when precise thermodynamic calculations are required. \n  It reflects how the properties of the solution deviate from ideal behavior and provides a more accurate understanding of the $pH$ under various conditions.\n  $$$$\n  While $pH$ is a universally recognized scale for expressing hydrogen ion activity,\n  its appropriateness and accuracy can diminish under conditions of extremely high\n  ionic strength, non-aqueous environments, high temperatures, or very high or low $pH$ values.\n  In such cases, alternative measurement strategies might be required to obtain meaningful and accurate descriptions of acidity or basicity.\n  " 
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    ,"label":"Acidity" 
    ,"symbol":"pH" 
    ,"type":"qudt:Unit" 
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JSON-LD

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  "description" : "\n  In chemistry the unit $\\textit{pH}$, also referred to as $\\textit{acidity}$ or $\\textit{basicity}$,\n   is the negative logarithm (base 10) of the concentration of free protons (or hydronium ions).\n  The definition of $pH$ in terms of hydrogen ions in solution is:\n\n  $$\\text{pH}=-\\log_{10}(a_{H^+})\\equiv-\\log_{10}\\left(\\left[H^+\\right]\\right)$$\n\n  Where $a_{H^+}$ is the equilibrium molar concentration of $H^+$ in the solution, the activity of\n   the hydrogen ion in the solution.\n  $$$$\n  This definition is appropriate for concentrations equal to, or less than $1\\ mol/l$, \n   where $aH+ \\equiv [H+]$, that is, $1\\ mol/L\\ HCl$ has a $pH$ of zero.\n  $$$$\n  To relate this to standard molality ($b^\\circ$), typically taken as $1 \\ mol/kg$, \n   consideration is given to the activity ($a_{H^+}$) of the hydrogen ions.\n  $$$$\n  The activity can be expressed as:\n\n  $$a_{H^+} = \\gamma_{H^+} \\times m_{H^+}$$ \n\n  Where, $\\gamma_{H^+}$ is the activity coefficient, which adjusts the molality to account for\n   non-ideal behavior due to interactions between ions in the solution.\n  $m_{H^+}$ is the molality of hydrogen ions in the solution relative to the standard molality, \n   expressed in $mol/kg$.\n  $$$$\n  The expansion of $pH$ then becomes:\n\n  $$\\text{pH} = -log_{10}\\left(m_{H+}\\times\\gamma_{H^+}\\right)$$\n\n  $$$$\n  This definition is relevant in more concentrated solutions or when precise thermodynamic calculations are required. \n  It reflects how the properties of the solution deviate from ideal behavior and provides a more accurate understanding of the $pH$ under various conditions.\n  $$$$\n  While $pH$ is a universally recognized scale for expressing hydrogen ion activity,\n  its appropriateness and accuracy can diminish under conditions of extremely high\n  ionic strength, non-aqueous environments, high temperatures, or very high or low $pH$ values.\n  In such cases, alternative measurement strategies might be required to obtain meaningful and accurate descriptions of acidity or basicity.\n  ",
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