|
#ALIH [[Larutan dapar]]
{{unreferenced|date=Februari 2014}}
{{rapikan}}
{{disambig info|Buffer|Buffer (disambiguasi)}}
{{asam dan basa}}
<!-- A '''buffer solution''' (more precisely, [[pH]] buffer or [[hydrogen ion]] buffer) is an [[aqueous solution]] consisting of a [[mixture]] of a [[weak acid]] and its [[conjugate base]], or vice versa. Its pH changes very little when a small or moderate amount of [[strong acid]] or [[Base (chemistry)#Strong bases|base]] is added to it and thus it is used to prevent changes in the pH of a solution. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. Many life forms thrive only in a relatively small pH range so they utilize a buffer solution to maintain a constant pH. In nature, the [[bicarbonate buffering system]] is used to regulate the pH of [[blood]].
==Principles of buffering==
[[File:Buffer titration.png|thumb|250px|left|Simulated [[titration]] of an acidified solution of a weak acid (pK<sub>a</sub> = 4.7) with alkali.]]
[[File:Buffer Wiki Edit.png|thumb|thunb|400px|Addition of [[hydroxide]] to an equilibrium mixture of a weak acid. HA, and its conjugate base, A<sup>-</sup> ]]
Buffer solutions achieve their resistance to pH change because of the presence of an equilibrium between the acid HA and its conjugate base A<sup>−</sup>.
:HA {{eqm}} H<sup>+</sup> + A<sup>−</sup>
When some [[strong acid]] is added to an equilibrium mixture of the weak acid and its [[conjugate base]], the equilibrium is shifted to the left, in accordance with [[Le Chatelier's principle]]. Because of this, the hydrogen ion concentration increases by less than the amount expected for the quantity of strong acid added.
Similarly, if strong alkali is added to the mixture the hydrogen ion concentration decreases by less than the amount expected for the quantity of alkali added. The effect is illustrated by the simulated titration of a weak acid with pK<sub>a</sub> = 4.7. The relative concentration of undissociated acid is shown in blue and of its conjugate base in red. The pH changes relatively slowly in the buffer region, pH = pK<sub>a</sub> ± 1, centered at pH = 4.7 where [HA] = [A<sup>−</sup>]. The hydrogen ion concentration decreases by less than the amount expected because most of the added hydroxide ion is consumed in the reaction
:OH<sup>−</sup> + HA → H<sub>2</sub>O + A<sup>−</sup>
and only a little is consumed in the neutralization reaction which results in an increase in pH.
:OH<sup>−</sup> + H<sup>+</sup> → H<sub>2</sub>O
Once the acid is more than 95% deprotonated the pH rises rapidly because most of the added alkali is consumed in the neutralization reaction.
===Buffer capacity===
Buffer capacity, β, is a quantitative measure of the resistance of a buffer solution to pH change on addition of hydroxide ions. It can be defined as follows.
:<math>\beta = \frac{dn}{d(p[H^+])}</math>
where dn is an infinitesimal amount of added base and d(p[H<sup>+</sup>]) is the resulting infinitesimal change in the [[cologarithm]] of the hydrogen ion concentration. With this definition the buffer capacity of a weak acid, with a dissociation constant K<sub>a</sub>, can be expressed as
:<math>\frac{dn}{d(pH)}=2.303\left(\frac{C_AK_a[H^+]}{\left(K_a+[H^+]\right)^2}\right)</math>
where C<sub>A</sub> is the analytical concentration of the acid.<ref>{{cite book|last1=Butler|first1=J.N.|title=Ionic Equilibrium: A Mathematical Approach|date=1964|publisher=Addison-Wesley|page=151}}</ref><ref name=Hulanicki/> pH is defined as -log<sub>10</sub>[H<sup>+</sup>]. The buffer capacity of a buffering agent is at a local maximum when p[H<sup>+</sup>] = pK<sub>a</sub>. It falls to 33% of the maximum value at p[H<sup>+</sup>] = pK<sub>a</sub> ± 1 and to 10% at p[H<sup>+</sup>] = pK<sub>a</sub> ± 1.5. For this reason the useful range is approximately pK<sub>a</sub> ± 1. Buffer capacity is proportional to the concentration of the buffering agent, C<sub>A</sub>, so dilute solutions have little buffer capacity.
[[File:Buffer1 12.png|thumb|200px|Buffer capacity for a 0.1 M solution of an acid with pK<sub>a</sub> of 7]]
Water itself is a buffering medium, even in the absence of an added buffering reagent. Its buffer capacity can be expressed as
:<math>\frac{dn}{d(pH)}=2.303\left([H^+]+[OH^-] \right)</math>
* At very low p[H<sup>+</sup>] the first term predominates and β increases in proportion to the hydrogen ion concentration; buffer capacity rises exponentially with pH.
* At very high p[H<sup>+</sup>] the second term predominates and β increases in proportion to the hydroxide ion concentration; buffer capacity rises exponentially with pH.
These properties are independent of the presence or absence of added buffering agents. They are concentration effects and reflect the fact that pH is related to the logarithm of the hydrogen ion concentration.
==Applications==
Buffer solutions are necessary to keep the correct pH for [[enzyme]]s in many organisms to work. Many enzymes work only under very precise conditions; if the pH moves outside of a narrow range, the enzymes slow or stop working and can [[denaturation (biochemistry)|denature]]. In many cases denaturation can permanently disable their catalytic activity.<ref name="Scorpio 2000">{{cite book |title=Fundamentals of Acids, Bases, Buffers & Their Application to Biochemical Systems |last=Scorpio |first=R. |year=2000 |publisher=|isbn=0-7872-7374-0}}</ref>
A buffer of [[carbonic acid]] (H<sub>2</sub>CO<sub>3</sub>) and [[bicarbonate]] (HCO<sub>3</sub><sup>−</sup>) is present in [[blood plasma]], to maintain a pH between 7.35 and 7.45.
Industrially, buffer solutions are used in [[fermentation (biochemistry)|fermentation]] processes and in setting the correct conditions for dyes used in colouring fabrics. They are also used in chemical analysis<ref name=Hulanicki>{{cite book |last= Hulanicki |first= A. |title= Reactions of acids and bases in analytical chemistry |publisher= Horwood |year= 1987 |isbn=0-85312-330-6}} (translation editor: Mary R. Masson)</ref> and calibration of pH meters.
The majority of biological samples that are used in research are made in buffers, especially [[phosphate buffered saline]] (PBS) at pH 7.4.
===Simple buffering agents===
:{| class="wikitable" style="text-align:center"
!Buffering agent!!pK<sub>a</sub>!!useful pH range
|-
|[[Citric acid]]||3.13, 4.76, 6.40||2.1–7.4
|-
|[[Acetic acid]]||4.8||3.8–5.8
|-
|[[potassium dihydrogenphosphate|KH<sub>2</sub>PO<sub>4</sub>]]||7.2|| 6.2–8.2
|-
|[[N-Cyclohexyl-2-aminoethanesulfonic acid|CHES]]||9.3|| 8.3–10.3
|-
|[[Borate]]||9.24||8.25–10.25
|}
For buffers in acid regions, the pH may be adjusted to a desired value by adding a strong acid such as [[hydrochloric acid]] to the buffering agent. For alkaline buffers, a strong base such as [[sodium hydroxide]] may be added. Alternatively, a buffer mixture can be made from a mixure of an acid and its conjugate base. For example, an acetate buffer can be made from a mixture of acetic acid and [[sodium acetate]]. Similarly an alkaline buffer can be made from a mixture of the base and its conjugate acid.
==="Universal" buffer mixtures===
By combining substances with p''K''<sub>a</sub> values differing by only two or less and adjusting the pH, a wide range of buffers can be obtained. [[Citric acid]] is a useful component of a buffer mixture because it has three p''K''<sub>a</sub> values, separated by less than two. The buffer range can be extended by adding other buffering agents.
The following mixtures ([[McIlvaine's buffer]] solutions) have a buffer range of pH 3 to 8.<ref>{{cite journal|last=McIlvaine|first=T.C.|year=1921|title=A buffer solution for colorimetric comparaison|journal=J. Biol. Chem.|volume=49|pages=183–186|url=http://www.jbc.org/content/49/1/183.full.pdf|issue=1}}</ref>
:{| class="wikitable" style="text-align:center"
! 0.2M Na<sub>2</sub>HPO<sub>4</sub> /mL
! 0.1M Citric Acid /mL
! pH...
|-
| 20.55
| 79.45
| style="background:#ff0000;" | 3.0
|-
| 38.55
| 61.45
| style="background:#ff7777;" |4.0
|-
| 51.50
| 48.50
| style="background:#ff7700;" | 5.0
|-
| 63.15
| 36.85
| style="background:#ffff00;" |6.0
|-
| 82.35
| 17.65
| style="background:#007777;" | 7.0
|-
| 97.25
| 2.75
|style="background:#0077ff;" | 8.0
|}
A mixture containing [[citric acid]], [[monopotassium phosphate]], [[boric acid]], and [[Barbital|diethyl barbituric acid]] can be made to cover the pH range 2.6 to 12.<ref>{{cite book |title=Vogel's textbook of quantitative chemical analysis |last=Mendham |first=J. |author2=Denny, R.C. |author3=Barnes, J.D. |author4=Thomas, M |edition=5th.|year=2000 |publisher=Pearson Education |location=Harlow |isbn=0-582-22628-7}} Appendix 5</ref>
Other universal buffers are Carmody buffer<ref name=carmody>{{cite journal|last=Carmody|first=Walter R.|title=Easily prepared wide range buffer series|journal=J. Chem. Educ.|year=1961|volume=38|issue=11|pages=559–560|doi=10.1021/ed038p559|url=http://dx.doi.org/10.1021/ed038p559|bibcode = 1961JChEd..38..559C }}</ref> and [[Britton-Robinson buffer]], developed in 1931.
===Common buffer compounds used in biology===
For effective range see [[#buffer capacity|Buffer capacity]], above.
{| class="wikitable" style="text-align:center"
|- bgcolor="#DDDD22"
! Common Name !!Structure !! [[Acid dissociation constant|pK<sub>a</sub>]]<br>at 25 °C !! Temp Effect<br>''d''pH/''d''T in (1/K)<ref>{{Cite web|title=Buffer Reference Center |url=http://www.sigmaaldrich.com/life-science/core-bioreagents/biological-buffers/learning-center/buffer-reference-center.html |publisher=Sigma-Aldrich | accessdate=2009-04-17}}</ref> !! Mol.<br>Weight
|-
| [[TAPS (buffer)|TAPS]]|| [[file:TAPS.svg|200px]] || 8.43 || −0.018 || 243.3
|-
| [[Bicine]] ||[[file:Bicine.png|150px]] || 8.35 || −0.018 || 163.2
|-
| [[Tris]] ||[[file:tris.png|100px]] || 8.06 || −0.028 || 121.14
|-
| [[Tricine]] ||[[file:Tricine.png|150px]] || 8.05 || −0.021 || 179.2
|-
| [[TAPSO (buffer)|TAPSO]] ||[[file:TAPSO.svg|200px]] ||7.635|| ||259.3
|-
| [[HEPES]] ||[[file: HEPES.png|200px]] || 7.48 || −0.014 || 238.3
|-
| [[TES (buffer)|TES]] ||[[file:TES free acid.svg|200px]] || 7.40 || −0.020 || 229.20
|-
| [[MOPS]] ||[[file:MOPS.png|150px]] || 7.20 || −0.015 || 209.3
|-
| [[PIPES]] || [[file:PIPES.svg|200px]]|| 6.76 || −0.008 || 302.4
|-
| [[Cacodylate]] ||[[file:Cacodylic acid.svg|100px]] || 6.27 || || 138.0
|-
| [[MES (buffer)|MES]] || [[file:MES.svg|150px]]|| 6.15 || −0.011 || 195.2
|-
|}
See also biological buffers:<ref>{{Cite web|title=Biological buffers |url=http://www.reachdevices.com/Protein/BiologicalBuffers.html |publisher=REACH Devices}}</ref>
== Calculating buffer pH ==
=== Monoprotic acids ===
First write down the equilibrium expression.
:HA {{eqm}} A<sup>−</sup> + H<sup>+</sup>
This shows that when the acid dissociates equal amounts of hydrogen ion and anion are produced. The equilibrium concentrations of these three components can be calculated in an [[ICE table]].
:{| class="wikitable" style="text-align:center"
|+ICE table for a monoprotic acid
|-
!width=50| ||width=50|[HA]||width=50|[A<sup>−</sup>]||width=50|[H<sup>+</sup>]
|-
|I||C<sub>0</sub>||0||y
|-
|C||-x||x||x
|-
|E||C<sub>0</sub>-x||x||x+y
|}
The first row, labelled I, lists the initial conditions: the concentration of acid is C<sub>0</sub>, initially undissociated, so the concentrations of A<sup>−</sup> and H<sup>+</sup> would be zero; y is the initial concentration of ''added'' strong acid, such as hydrochloric acid. If strong alkali, such as sodium hydroxide, is added y will have a negative sign because alkali removes hydrogen ions from the solution. The second row, labelled C for change, specifies the changes that occur when the acid dissociates. The acid concentration decreases by an amount ''-x'' and the concentrations of A<sup>−</sup> and H<sup>+</sup> both increase by an amount ''+x''. This follows from the equilibrium expression. The third row, labelled E for equilibrium concentrations, adds together the first two rows and shows the concentrations at equilibrium.
To find ''x'', use the formula for the equilibrium constant in terms of concentrations:
:<math>K_a = \frac{[H^+] [A^-]}{[HA]}</math>
Substitute the concentrations with the values found in the last row of the ICE table:
:<math>K_a = \frac{x(x+y)}{C_0 - x}</math>
Simplify to:
:<math>x^2 + (K_a +y) x - K_a C_0 = 0</math>
With specific values for C<sub>0</sub>, K<sub>a</sub> and y this equation can be solved for x. Assuming that pH = -log<sub>10</sub>[H<sup>+</sup>] the pH can be calculated as pH = -log<sub>10</sub>(x+y).
===Polyprotic acids===
[[File:Citric acid speciation.png|thumb|200 px|alt=This image plots the relative percentages of the protonation species of citric acid as a function of p H. Citric acid has three ionizable hydrogen atoms and thus three p K A values. Below the lowest p K A, the triply protonated species prevails; between the lowest and middle p K A, the doubly protonated form prevails; between the middle and highest p K A, the singly protonated form prevails; and above the highest p K A, the unprotonated form of citric acid is predominant.| [[Determination of equilibrium constants#speciation calculations|% species formation]] calculated for a 10 millimolar solution of citric acid.]]
Polyprotic acids are acids that can lose more than one proton. The constant for dissociation of the first proton may be denoted as ''K''<sub>a1</sub> and the constants for dissociation of successive protons as ''K''<sub>a2</sub>, etc. [[Citric acid]], H<sub>3</sub>A, is an example of a polyprotic acid as it can lose three protons.
:{| class="wikitable"
!equilibrium!!p''K''<sub>a</sub> value
|-
| H<sub>3</sub>A {{eqm}} H<sub>2</sub>A<sup>−</sup> + H<sup>+</sup>
| p''K''<sub>a1</sub> = 3.13
|-
| H<sub>2</sub>A<sup>−</sup> {{eqm}} HA<sup>2−</sup> + H<sup>+</sup>
| p''K''<sub>a2</sub> = 4.76
|-
| HA<sup>2−</sup> {{eqm}} A<sup>3−</sup> + H<sup>+</sup>
| p''K''<sub>a3</sub> = 6.40
|}
When the difference between successive p''K'' values is less than about three there is overlap between the pH range of existence of the species in equilibrium. The smaller the difference, the more the overlap. In the case of citric acid, the overlap is extensive and solutions of citric acid are buffered over the whole range of pH 2.5 to 7.5.
Calculation of the pH with a polyprotic acid requires a [[Determination of equilibrium constants#Speciation calculations|speciation calculation]] to be performed. In the case of citric acid, this entails the solution of the two equations of mass balance
:<math> C_A = [A^{3-}]+\beta_1 [A^{3-}][H^+] +\beta_2 [A^{3-}][H^+]^2 +\beta_3 [A^{3-}][H^+]^3</math>
:<math> C_H = [H^+]+ \beta_1 [A^{3-}][H^+]+ 2\beta_2 [A^{3-}][H^+]^2+ 3\beta_3 [A^{3-}][H^+]^3 -K_w[H^+]^{-1}</math>
C<sub>A</sub> is the analytical concentration of the acid, C<sub>H</sub> is the analytical concentration of added hydrogen ions, β<sub>q</sub> are the [[equilibrium constant#Cumulative and stepwise formation constants|cumulative association constants]]
:<math>\log \beta_1=pK_{a3}, \ \log \beta_2=pK_{a2}+ pK_{a3},\ \log \beta_3=pK_{a1}+ pK_{a2}+ pK_{a3} </math>
K<sub>w</sub> is the constant for [[Self-ionization of water]]. There are two [[non-linear]] [[simultaneous equation]]s in two unknown quantities [A<sup>3−</sup>] and [H<sup>+</sup>]. Many computer programs are available to do this calculation. The speciation diagram for citric acid was produced with the program HySS.<ref>{{cite journal | last1 = Alderighi | first1 = L. | last2 =Gans | first2 = P. | last3 = Ienco | first3 = A. | last4 = Peters | first4 = D. | last5 = Sabatini | first5 = A. | last6 = Vacca| first6 = A. | year = 1999 | title = Hyperquad simulation and speciation (HySS): a utility program for the investigation of equilibria involving soluble and partially soluble species | journal = Coordination Chemistry Reviews | volume = 184 | issue = 1 | pages = 311–318 | doi = 10.1016/S0010-8545(98)00260-4 | url = http://www.hyperquad.co.uk/hyss.htm}}</ref>
==See also==
* [[Henderson–Hasselbalch equation]]
* [[Buffering agent]]
* [[Good's buffers]]
* [[Common-ion effect]]
* [[Metal ion buffer]]
* [[Mineral redox buffer]]
==References==
{{reflist}}
==External links==
* [http://www.popproperty.net/PopularTools/PHBuffer1.aspx Online pH buffer calculator]
* [http://www.cnr.berkeley.edu/soilmicro/methods/phosphate%20buffer.pdf phosphate buffer]
{{Chemical equilibria}}
[[Category:Acid–base chemistry]]
[[Category:Equilibrium chemistry]]
[[Category:Buffers]]
-->'''Larutan penyangga''', '''larutan dapar''', atau '''''buffer''''' adalah [[larutan]] yang digunakan untuk mempertahankan nilai [[pH]] tertentu agar tidak banyak berubah selama [[reaksi kimia]] berlangsung. Sifat yang khas dari larutan penyangga ini adalah pH-nya hanya berubah sedikit dengan pemberian sedikit asam kuat atau basa kuat. Sehingga pH akhirnya tidak jauh berbeda dengan pH awal.
Larutan penyangga tersusun dari [[asam]] lemah dengan [[basa]] [[konjugat]]nya atau oleh basa lemah dengan asam konjugatnya. Reaksi di antara kedua komponen penyusun ini disebut sebagai reaksi asam-basa konjugasi.
== Komponen Larutan Penyangga ==
'''Secara umum, larutan penyangga digambarkan sebagai campuran yang terdiri dari:'''
*Asam lemah (HA) dan basa konjugasinya (ion A-), campuran ini menghasilkan larutan bersifat asam.
*Basa lemah (B) dan asam konjugasinya (BH+), campuran ini menghasilkan larutan bersifat basa.
'''Komponen larutan penyangga terbagi menjadi:'''
*Larutan penyangga yang bersifat asam
Larutan ini mempertahankan pH pada daerah asam (pH < 7). Untuk mendapatkan larutan ini dapat
dibuat dari asam lemah dan garamnya yang merupakan basa konjugasi dari asamnya. Adapun cara lainnya yaitu mencampurkan suatu asam lemah dengan suatu basa kuat di mana asam lemahnya dicampurkan dalam jumlah berlebih. Campuran akan menghasilkan garam yang mengandung basa konjugasi dari asam lemah yang bersangkutan. Pada umumnya basa kuat yang digunakan seperti natriumNa), kalium, barium, kalsium, dan lain-lain.
*Larutan penyangga yang bersifat basa
Larutan ini mempertahankan pH pada daerah basa (pH > 7). Untuk mendapatkan larutan ini dapat dibuat dari basa lemah dan garam, yang garamnya berasal dari asam kuat. Adapun cara lainnya yaitu dengan mencampurkan suatu basa lemah dengan suatu asam kuat di mana basa lemahnya dicampurkan berlebih.
== Cara kerja larutan penyangga ==
Larutan penyangga mengandung komponen asam dan basa dengan asam dan basa konjugasinya, sehingga dapat mengikat baik ion H+ maupun ion OH-. Sehingga penambahan sedikit asam kuat atau basa kuat tidak mengubah pH-nya secara signifikan. Berikut ini cara kerja larutan penyangga:
===Larutan penyangga asam===
Adapun cara kerjanya dapat dilihat pada larutan penyangga yang mengandung CH3COOH dan CH3COO- yang mengalami kesetimbangan. Dengan proses sebagai berikut:
*'''Pada penambahan asam'''
Penambahan [[asam]] (H+) akan menggeser kesetimbangan ke kiri. Di mana ion H+ yang ditambahkan akan bereaksi dengan ion CH3COO- membentuk molekul CH3COOH.
<center>'''CH3COO-(aq) + H+(aq) → CH3COOH(aq''')</center>
*'''Pada penambahan basa'''
Jika yang ditambahkan adalah suatu basa, maka ion OH- dari basa itu akan bereaksi dengan ion H+ membentuk air. Hal ini akan menyebabkan kesetimbangan bergeser ke kanan sehingga konsentrasi ion H+ dapat dipertahankan. Jadi, penambahan basa menyebabkan berkurangnya komponen asam (CH3COOH), bukan ion H+. Basa yang ditambahkan tersebut bereaksi dengan asam CH3COOH membentuk ion CH3COO- dan air.
<center>'''CH3COOH(aq) + OH-(aq) → CH3COO-(aq) + H2O(l)'''</center>
===Larutan penyangga basa===
Adapun cara kerjanya dapat dilihat pada larutan penyangga yang mengandung NH3 dan NH4+ yang mengalami kesetimbangan. Dengan proses sebagai berikut:
*'''Pada penambahan asam'''
Jika ditambahkan suatu asam, maka ion H+ dari asam akan mengikat ion OH-. Hal tersebut menyebabkan kesetimbangan bergeser ke kanan, sehingga konsentrasi ion OH- dapat dipertahankan. Disamping itu penambahan ini menyebabkan berkurangnya komponen basa (NH3), bukannya ion OH-. Asam yang ditambahkan bereaksi dengan basa NH3 membentuk ion [[NH4+]].
<center>'''NH3 (aq) + H+(aq) → NH4+ (aq)'''</center>
*'''Pada penambahan basa'''
Jika yang ditambahkan adalah suatu basa, maka kesetimbangan bergeser ke kiri, sehingga konsentrasi ion OH- dapat dipertahankan. Basa yang ditambahkan itu bereaksi dengan komponen asam (NH4+), membentuk komponen basa (NH3) dan air.
<center>'''NH4+ (aq) + OH-(aq) → NH3 (aq) + H2O(l)'''</center>
== Perhitungan pH Larutan Penyangga ==
===Larutan penyangga asam===
Dapat digunakan tetapan ionisasi dalam menentukan konsentrasi ion H+ dalam suatu larutan dengan rumus berikut:
<center>'''[H+] = Ka x a/valxg'''</center>
<center>'''atau'''</center>
<center>'''pH = p Ka - log a/g'''</center>
dengan, Ka = tetapan ionisasi asam lemah
::a = jumlah mol asam lemah
::g = jumlah mol basa konjugasi
===Larutan penyangga basa===
Dapat digunakan tetapan ionisasi dalam menentukan konsentrasi ion H+ dalam suatu larutan dengan rumus berikut:
<center>'''[OH-] = Kb x b/valxg'''</center>
<center>'''''atau'''''</center>
<center>'''pOH = p Kb - log b/g'''
'''pH = 14 - pOH'''
</center>
dengan, Kb = tetapan ionisasi basa lemah
::b = konsentrasi basa lemah
::g = konsentrasi asam konjugasi
== Fungsi Larutan Penyangga ==
Adanya larutan penyangga ini dapat kita lihat dalam kehidupan sehari-hari seperti pada obat-obatan, fotografi, industri kulit dan zat warna. Selain aplikasi tersebut, terdapat fungsi penerapan konsep larutan penyangga ini dalam tubuh manusia seperti pada cairan tubuh.
Cairan tubuh ini bisa dalam cairan intrasel maupun cairan ekstrasel. Di mana sistem penyangga utama dalam cairan intraselnya seperti H2PO4- dan HPO42- yang dapat bereaksi dengan suatu asam dan basa. Adapun sistem penyangga tersebut, dapat menjaga pH darah yang hampir konstan yaitu sekitar 7,4.
Selain itu penerapan larutan penyangga ini dapat kita temui dalam kehidupan sehari-hari seperti pada obat tetes mata. Pada obat tetes mata mempunyai pH yang sama dengan cairan tubuh kita, agar tidak menimbulkan efek samping.
== Pranala luar ==
* [http://www.popproperty.net/PopularTools/PHBuffer1.aspx Online pH buffer calculator]
* [http://www.cnr.berkeley.edu/soilmicro/methods/phosphate%20buffer.pdf phosphate buffer]
[[Kategori:Kimia|Larutan]]
| 