Bronsted-Lowry acid and bases
A Bronsted-Lowry acid (proton donor) is a chemical substance that is able to donate protons or a H+
ions to another substance during a reaction, for example CH3COOH + H2O ⇌ H3O + + CH3COO−. A
Bronsted-Lowry base (proton acceptor) is also a chemical substance that accepts protons and H+ ions
from another substance during a reaction, for example C6H 5NH 2+ H2O ⇌ C6H5NH3 + + OH−. A
common Bronsted-Lowry acid-base reaction is between hydrochloric acid and a hydroxide ion shown
below.
During this rection, a proton is transferred from hydrochloric acid (proton donor) to the hydroxide
ion (proton acceptor). Once the Bronsted-Lowry acid donated a proton, a conjugate base remains,
this is a species that may take a proton or hydrogen ion. It is created by taking a proton out of an
acid. A pair of chemicals that can change into one another by receiving or losing a proton are said to
be conjugate bases. In other words, a conjugate acid or base are two species that differ by H +, for
example an acid and its conjugate base, such as H3O + and H2O or a base and its conjugate acid,
such as NH3 , NH4 +. Acetate, for instance, is acetic acid's conjugate base and can receive a proton to
change back into acetic acid. The chloride ion then becomes the conjugate base of hydrochloric acid.
pH Scale
pH is a figure expressing the acidity or alkalinity of a solution on a logarithmic scale on which 7 is
neutral, lower values are more acidic and higher values more alkaline. Each substance also has a
concentration of [H+] and [OH-], a pH value, and a pOH value. All bases fall below 7 on the pOH
scale. A strong base would be in the range of 1 to 3, whereas a weak base would be in the range of 4
to 6. Above 7, which is still neutral, is an acid. Between 8 and 10, there would be a weak acid, and
between 11 and 14, a strong acid. The concentration of bases is directly connected to pOH and can
be used to calculate it, just like it is for acids. Concentration in bases, denoted as [OH-], is the
concentration of OH- rather than H+. These can be solved using the equations below.
, Calculating pH from [H+]
pH= -log[H+]
Calculating pOH from [OH-]
pOH= -log[OH-]
Calculating pOH from pH
pOH=14-pH
Calculating pH from pOH
pH=14-pOH
Calculating [H+] from pH
[H+]= inverse log(-pH)
which is the same as
[H+]= 10^-pH
Calculating [OH-] from pOH
[OH-]= inverse log(-pOH)
which is the same as
[OH-]= 10^-pOH
pH = -log [H+] or. pH equals log1/[H+] is used to calculate the pH of an aqueous solution is the
negative logarithm of the hydrogen ion concentration (in mol per litre) or pH of the solution is the
logarithm of the reciprocal of H+ ion concentration. [H+] is the acid concentration for strong acids.
Degree of dissociation
The relative strength of an acid is determined by how quickly it ionises in water. The Ka is only the
ionisation of an acid, HA, into H+ and A- equilibrium constant. The expression Ka [H+][A-]/[HA] can
be written. Using the above-mentioned formula, which includes concentration of the solution and
Ka, which can be readily determine the degree of dissociation to be. It can be concluded that
stronger acids resemble those with higher Ka values. A lower pKa value with a negative logKa will
therefore resemble a stronger acid.
Ka
Ka is the acidic dissociation constant, Ka = ([H + ][A − ])/[HA]. Taking the negative logarithm and
rearranging for pH gives: pH = pKa + log([A − ]/[HA]). Ka or acid dissociation constant is the
quantitative measure of the strength of an acid. It tells us about how much an acid dissociates in an
aqueous solution. The larger the Ka value, the more would be its dissociation in water. It is used to
distinguish strong acids from weak acids.
pKa
A Bronsted-Lowry acid (proton donor) is a chemical substance that is able to donate protons or a H+
ions to another substance during a reaction, for example CH3COOH + H2O ⇌ H3O + + CH3COO−. A
Bronsted-Lowry base (proton acceptor) is also a chemical substance that accepts protons and H+ ions
from another substance during a reaction, for example C6H 5NH 2+ H2O ⇌ C6H5NH3 + + OH−. A
common Bronsted-Lowry acid-base reaction is between hydrochloric acid and a hydroxide ion shown
below.
During this rection, a proton is transferred from hydrochloric acid (proton donor) to the hydroxide
ion (proton acceptor). Once the Bronsted-Lowry acid donated a proton, a conjugate base remains,
this is a species that may take a proton or hydrogen ion. It is created by taking a proton out of an
acid. A pair of chemicals that can change into one another by receiving or losing a proton are said to
be conjugate bases. In other words, a conjugate acid or base are two species that differ by H +, for
example an acid and its conjugate base, such as H3O + and H2O or a base and its conjugate acid,
such as NH3 , NH4 +. Acetate, for instance, is acetic acid's conjugate base and can receive a proton to
change back into acetic acid. The chloride ion then becomes the conjugate base of hydrochloric acid.
pH Scale
pH is a figure expressing the acidity or alkalinity of a solution on a logarithmic scale on which 7 is
neutral, lower values are more acidic and higher values more alkaline. Each substance also has a
concentration of [H+] and [OH-], a pH value, and a pOH value. All bases fall below 7 on the pOH
scale. A strong base would be in the range of 1 to 3, whereas a weak base would be in the range of 4
to 6. Above 7, which is still neutral, is an acid. Between 8 and 10, there would be a weak acid, and
between 11 and 14, a strong acid. The concentration of bases is directly connected to pOH and can
be used to calculate it, just like it is for acids. Concentration in bases, denoted as [OH-], is the
concentration of OH- rather than H+. These can be solved using the equations below.
, Calculating pH from [H+]
pH= -log[H+]
Calculating pOH from [OH-]
pOH= -log[OH-]
Calculating pOH from pH
pOH=14-pH
Calculating pH from pOH
pH=14-pOH
Calculating [H+] from pH
[H+]= inverse log(-pH)
which is the same as
[H+]= 10^-pH
Calculating [OH-] from pOH
[OH-]= inverse log(-pOH)
which is the same as
[OH-]= 10^-pOH
pH = -log [H+] or. pH equals log1/[H+] is used to calculate the pH of an aqueous solution is the
negative logarithm of the hydrogen ion concentration (in mol per litre) or pH of the solution is the
logarithm of the reciprocal of H+ ion concentration. [H+] is the acid concentration for strong acids.
Degree of dissociation
The relative strength of an acid is determined by how quickly it ionises in water. The Ka is only the
ionisation of an acid, HA, into H+ and A- equilibrium constant. The expression Ka [H+][A-]/[HA] can
be written. Using the above-mentioned formula, which includes concentration of the solution and
Ka, which can be readily determine the degree of dissociation to be. It can be concluded that
stronger acids resemble those with higher Ka values. A lower pKa value with a negative logKa will
therefore resemble a stronger acid.
Ka
Ka is the acidic dissociation constant, Ka = ([H + ][A − ])/[HA]. Taking the negative logarithm and
rearranging for pH gives: pH = pKa + log([A − ]/[HA]). Ka or acid dissociation constant is the
quantitative measure of the strength of an acid. It tells us about how much an acid dissociates in an
aqueous solution. The larger the Ka value, the more would be its dissociation in water. It is used to
distinguish strong acids from weak acids.
pKa
