Biotecnikaglobal

ONE STOP SITE FOR ALL YOUR BIOTECHNOLOGY NEEDS

CHEMICAL FOUNDATIONS OF BIOLOGY

pH

pH is a measure of the acidity of a solution in terms of activity of hydrogen (H⁺). For dilute solutions, however, it is convenient to substitute the activity of the hydrogen ions with the molarity (mol/L) of the hydrogen ions (however, this is not necessarily accurate at higher concentrations

In aqueous systems, the hydrogen ion activity is dictated by the dissociation constant of water (K_w = 1.011 × 10⁻¹⁴ M² at 25 °C) and interactions with other ions in solution. Due to this dissociation constant, a neutral solution (hydrogen ion activity equals hydroxide ion activity) has a pH of approximately 7. Aqueous solutions with pH values lower than 7 are considered acidic, while pH values higher than 7 are considered basic.

The concept was introduced by S.P.L. Sørensen in 1909, and is purported to mean pondus hydrogenii in Latin. However, most other sources attribute the name to the French term pouvoir hydrogène. In English, pH can stand for "hydrogen power,"power of hydrogen," or "potential of hydrogen.All of these terms are technically correct.

Acid

An acid (often represented by the generic formula HA) is traditionally considered any chemical compound that, when dissolved in water, gives a solution with a pH of less than 7.0. That approximates the modern definition of Brønsted and Lowry, who defined an acid as a compound which donates a hydrogen ion (H⁺) to another compound (called a base). Common examples include acetic acid (in vinegar) and sulfuric acid (used in car batteries).

Properties

Generally, acids have the following properties:

Taste: Acids generally are sour when dissolved in water.
Touch: Acids produce a stinging feeling, particularly strong acids.
Reactivity: Acids react aggressively with or corrode most metals.
Electrical conductivity: Acids, while not normally ionic, are electrolytes.

Strong acids

Strong acid and most concentrated acids are dangerous, causing severe burns for even minor contact. Generally, acid burns are treated by rinsing the affected area abundantly with running water (15 minutes) and followed up with immediate medical attention. In

the case of highly concentrated acids, the acid should first be wiped off as much as possible, otherwise the reaction of the acid dissolving in the water could cause severe thermal burns. In addition to dangers from the acidity, even dilute solutions of weak acids may also be dangerous, due to toxic or other effects of the ions involved.

Weak acid

In order to lose a proton, it is necessary that the pH of the system rise above the pK_a of the protonated acid. The decreased concentration of H⁺ in that basic solution shifts the equilibrium towards the conjugate base form (the deprotonated form of the acid). In lower-pH (more acidic) solutions, there is a high enough H⁺ concentration in the solution to cause the acid to remain in its protonated form, or to protonate its conjugate base (the deprotonated form).

Base

In chemistry, a base is most commonly thought of as a substance which can accept protons. This refers to the Bronsted-Lowry theory of acids and bases. Alternate definitions of bases include electron pair donors (Lewis), and as sources of hydroxide anions (Arrhenius). Examples of simple bases are sodium hydroxide and ammonia.

Bases can be thought of as the chemical opposite of acids. A reaction between an acid and base is called neutralization. Bases and acids are seen as opposites because the effect of an acid is to increase the hydronium ion (H₃O⁺) concentration in water, where as bases reduce this concentration. Bases react with acids to produce water and salts (or their solutions). Some general properties of bases include:

Taste: Bitter taste (opposed to sour taste of acids and sweetness of aldehydes and ketones)
Touch: Slimy or soapy feel on fingers
Reactivity: Caustic on organic matter, react violently with acidic or reducible substances
Electric conductivity: Aqueous solutions or molten bases dissociate in ions and conduct electricity
Bases turn red litmus paper blue

The Proton in Water: Arrhenius Theory

Swedish chemist Arrhenius defined an acid as a substance that ionizes in water to give hydrogen ions, and a base as a substance that ionizes in water to give hydroxide ions.

Hydrochloric acid, HCl, is a strong acid, and is very soluble in water. It dissociates into its component ions in the following manner:

HCl (g) ==>H⁺(aq) + Cl^-(aq)

The hydrogen ion interacts strongly with a lone pair of electrons on the oxygen of a water molecule. The resulting ion, H₃O⁺ is called the hydronium ion.

H⁺ + H₂O ==> H₃0⁺

ACIDIC solutions are formed when an acid transfers a proton to water.

The reaction of HCl with water can be written in either of the following ways:

HCl (aq) + H₂O (l) H₃O⁺ (aq) + Cl^-(aq)

HCl (aq) H⁺ (aq) + Cl^- (aq)

The Brønsted-Lowry Concept of Acids and Bases

Acids are substances that are capable of donating a proton, and bases are substances capable of accepting a proton.

So, in the example above, HCl acts as a Brønsted Acid by donating a proton in water, and water in turn acts as a Brønsted Base by accepting a proton from HCl. (as shown in this animation).

Water can act as an acid or a base. Here is another example:

NH₃(aq) + H₂O(l) NH₄⁺(aq) + OH^-(aq)

Here, H₂O acts as a Brønsted acid by donating a proton to NH₃ which acts as a Brønsted base.
Using the Arrhenius definition, we say that the resulting solution is basic because it contains OH^- ions, thus we say that the NH₃ molecule is basic (a proton acceptor).

All Arrhenius acids are also Brønsted acids.
All Arrhenius bases are also Brønsted bases.

Conjugate Acid-Base Pairs

Let's look at the reaction of NH₃ and H₂O again:

(1) NH₃ + H₂O NH₄⁺ + OH^-

The reverse of this reaction is:

(2) NH₄ + OH^- NH₃ + H₂O

In this case, NH₄⁺ acts as an acid which donates a proton to OH^-. OH^- acts as a base.

An acid and a base that are related by the gain and loss of a proton are called a conjugate acid-base pair. For example, NH₄⁺ is the conjugate acid of NH₃, and NH₃ is the conjugate base of NH₄⁺.

Every acid has associated with it a conjugate base.
Likewise, every base has associated with it a conjugate acid.

NH₃(aq) + H₂0 (l) NH₄⁺ (aq) + OH^-(aq)

NH₄⁺ (aq) + OH^-(aq) NH₃ (aq) + H₂O (l)

Acids and bases can be either neutral or charged species: H₂O (acid or base, neutral), O₂^- (base, charged), C₂H₃O₂^- (acid, charged)

For any reaction:

HA + H₂O H₃O⁺ + A^-

If HA is a strong acid because it gives up its proton readily, then A^- is a weak base because it has little affinity for the proton.

If HA is a weak acid because it donates very few protons to the water, then A^- has a high affinity for a proton, and A^- is a stronger base than water.

Henderson-Hasselbalch Equation

An acid is a substance which produces hydrogen ions (H⁺) by dissociation.
For example HCl ---> H⁺ + Cl^-

Bases are substances which can combine with H⁺ like ammonia (NH₃)
For example H⁺ + NH₃ --> NH₄⁺

They may produce hydroxide ions (OH^- ) either by direct dissociation or subsequent to reaction with water,
e.g. KOH ---> K⁺ + OH^-
and NH₃ + H₂O NH₄OH NH₄⁺ + OH^-

Hydrogen ions (H⁺) associate with water to form H₃O⁺ (and H₁₁O₅⁺ etc.)

Water dissociates to a tiny extent: H₂O H⁺ + OH^-
    The equilibrium constant = K_eq
                                               = 1.8 x 10^-16 = [H⁺]x[OH^-] = [H⁺]x[OH⁺]
                                                                           [H₂O]          (1000/18)

Therefore [H⁺]x[OH^- ] = 1.8 x 10-16 x 55.5 = 10^-14

In pure water [H⁺] must equal [OH^-] (the solution is neutral),
therefore [H⁺] = Ö10^-14 = 10^-7 M

The pH (hydrogen ion potential) of a solution is defined as pH = -log₁₀ (H⁺) , where (H⁺) is the hydrogen ion concentration. The pH scale can range between about -1 and +15; a neutral solution has pH 7.0.

For a weak acid, which dissociates as follows: HA H⁺ + A^-