unit:PH

URI: http://qudt.org/vocab/unit/PH

Type
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 ($b^\circ$), consideration is given to the activity ($a_{H^+}$) of the hydrogen ions is the product of their molality ($m_{H^+}$), and the activity coefficient ($\gamma_{H^+}$), where $m_{H^+}$ is expressed in $mol/kg$, and $b^\circ$ is typically taken as $1 \ mol/kg$. The activity can then 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. The expansion of $pH$ then becomes: $$\text{pH} = -log_{10}\left(m_{H+}\times\gamma_{H^+}\right)$$ Where, $m_{H+}$ is the molality of hydrogen ions in the solution relative to the standard molality, which can also be expressed as $b^{\circ}$. 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.

Properties
qudt:applicableSystem
qudt:hasDimensionVector
qudt:hasQuantityKind
qudt:symbol
pH
rdf:type
qudt:ucumCode
[pH]
Annotations
rdfs:comment
Unsure about dimensionality of pH; conversion requires a log function not just a multiplier(en)
dcterms: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 ($b^\circ$), consideration is given to the activity ($a_{H^+}$) of the hydrogen ions is the product of their molality ($m_{H^+}$), and the activity coefficient ($\gamma_{H^+}$), where $m_{H^+}$ is expressed in $mol/kg$, and $b^\circ$ is typically taken as $1 \ mol/kg$. The activity can then 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. The expansion of $pH$ then becomes: $$\text{pH} = -log_{10}\left(m_{H+}\times\gamma_{H^+}\right)$$ Where, $m_{H+}$ is the molality of hydrogen ions in the solution relative to the standard molality, which can also be expressed as $b^{\circ}$. 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.
qudt:informativeReference
https://iupac.org/wp-content/uploads/2019/05/IUPAC-GB3-2012-2ndPrinting-PDFsearchable.pdf
rdfs:isDefinedBy
rdfs:label
Acidity(en)
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 ($b^\circ$), consideration is given to the activity ($a_{H^+}$)
   of the hydrogen ions is the product of their molality ($m_{H^+}$), and the activity coefficient ($\gamma_{H^+}$),
   where $m_{H^+}$ is expressed in $mol/kg$, and $b^\circ$ is typically taken as $1 \ mol/kg$.
  
  
  The activity can then 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.

  The expansion of $pH$ then becomes:

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

  Where, $m_{H+}$ is the molality of hydrogen ions in the solution relative to the standard 
    molality, which can also be expressed as $b^{\circ}$.

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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<http://qudt.org/vocab/unit/PH>
  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 ($b^\\circ$), consideration is given to the activity ($a_{H^+}$)
   of the hydrogen ions is the product of their molality ($m_{H^+}$), and the activity coefficient ($\\gamma_{H^+}$),
   where $m_{H^+}$ is expressed in $mol/kg$, and $b^\\circ$ is typically taken as $1 \\ mol/kg$.
  
  
  The activity can then 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.

  The expansion of $pH$ then becomes:

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

  Where, $m_{H+}$ is the molality of hydrogen ions in the solution relative to the standard 
    molality, which can also be expressed as $b^{\\circ}$.

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:IMPERIAL" 
    ,"applicable system":"sou:PLANCK" 
    ,"applicable system":"sou:SI" 
    ,"comment":"Unsure about dimensionality of pH; conversion requires a log function not just a multiplier" 
    ,"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 ($b^\\circ$), consideration is given to the activity ($a_{H^+}$)\n   of the hydrogen ions is the product of their molality ($m_{H^+}$), and the activity coefficient ($\\gamma_{H^+}$),\n   where $m_{H^+}$ is expressed in $mol\/kg$, and $b^\\circ$ is typically taken as $1 \\ mol\/kg$.\n  \n  \n  The activity can then 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\n  The expansion of $pH$ then becomes:\n\n  $$\\text{pH} = -log_{10}\\left(m_{H+}\\times\\gamma_{H^+}\\right)$$\n\n  Where, $m_{H+}$ is the molality of hydrogen ions in the solution relative to the standard \n    molality, which can also be expressed as $b^{\\circ}$.\n\nThis definition is relevant in more concentrated solutions or when precise thermodynamic\n calculations are required. \nIt reflects how the properties of the solution deviate from ideal behavior and provides a more accurate\n understanding of the $pH$ under various conditions.\n\nWhile $pH$ 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\n  low $pH$ values.\nIn such cases, alternative measurement strategies might be required to obtain\n   meaningful and accurate descriptions of acidity or basicity.\n" 
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    ,"has quantity kind":"quantitykind:Basicity" 
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    ,"isDefinedBy":"&lt;http:\/\/qudt.org\/2.1\/vocab\/unit&gt;" 
    ,"label":"Acidity" 
    ,"symbol":"pH" 
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    ]}
JSON-LD 
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  "@type" : "http://qudt.org/schema/qudt/Unit",
  "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 ($b^\\circ$), consideration is given to the activity ($a_{H^+}$)\n   of the hydrogen ions is the product of their molality ($m_{H^+}$), and the activity coefficient ($\\gamma_{H^+}$),\n   where $m_{H^+}$ is expressed in $mol/kg$, and $b^\\circ$ is typically taken as $1 \\ mol/kg$.\n  \n  \n  The activity can then 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\n  The expansion of $pH$ then becomes:\n\n  $$\\text{pH} = -log_{10}\\left(m_{H+}\\times\\gamma_{H^+}\\right)$$\n\n  Where, $m_{H+}$ is the molality of hydrogen ions in the solution relative to the standard \n    molality, which can also be expressed as $b^{\\circ}$.\n\nThis definition is relevant in more concentrated solutions or when precise thermodynamic\n calculations are required. \nIt reflects how the properties of the solution deviate from ideal behavior and provides a more accurate\n understanding of the $pH$ under various conditions.\n\nWhile $pH$ 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\n  low $pH$ values.\nIn such cases, alternative measurement strategies might be required to obtain\n   meaningful and accurate descriptions of acidity or basicity.\n",
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