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Diagram of formic acid

The word "acid" comes from the Latin acidus meaning "sour," but in chemistry the term acid has a more specific meaning, which is most commonly accepted to be a chemical substance that acts as a proton donor (H+).


Acids have many properties that stand out. Some that you might be familiar to is in food and household cleaning products. It gives your food the sour taste and cleaning products the power to kill germs and erode stains. Acids react with compounds containing hydroxide ions to form water and a salt. There are several main types of acids: Citric, Malic, Lactic, Acidic, succinic, and Tartaric. Tartaric and malic acid are found naturally in grape juices. Citric is a less common acid found in that. Acetic and succinic are found in wines/must along with Lactic, and Malic. Tartaric Acid is the main acid added to increase the acidity or lower the pH of grape juice or wine. Citric Acid is often added for acidification as it is claimed to taste fruitier. [1]


There are four definitions of acids:

The Arrhenius
"This acid is a substance that increases the concentration of hydrogen ions (H+), which are carried as hydronium ions (H3O+</sup>) when dissolved in water, while bases are substances that increase the concentration of hydroxide ions (OH-). This definition limits acids and bases to substances that can dissolve in water."

The Brønsted-Lowry
"This acid is a proton (hydrogen nucleus) donor and a base is a proton acceptor. The acid is said to be dissociated after the proton is donated. An acid and the corresponding base are referred to as conjugate acid-base pairs. Brønsted and Lowry independently formulated this definition, which includes water-insoluble substances not in the Arrhenius definition."

"This acid is an electron-pair acceptor and a base is an electron-pair donor. Lewis acids include substances with no transferable protons."

"An acid is a substance that, when dissolved in an auto dissociating solvent, increases the concentration of the solvonium cations, such as H3O+ in water, NH4+ in liquid ammonia, NO+ in liquid N2O4, SbCl2+ in SbCl3, etc." [2]

Common acid reactions

Main Article: Acid–base reaction
lemon is very acidic with a pH of 4

Acids are generally sour in taste. Strong or concentrated acids often produce a stinging feeling on mucous membranes. React to indicators as follows: turn blue litmus and methyl orange red, do not change the color of phenolphthalein

  • Will react with metals to produce a metal salt and hydrogen
  • Will react with metal carbonates to produce water, CO2 and a salt
  • Will react with a base to produce a salt and water
  • Will react with a metal oxide to produce water and a salt
  • Will conduct electricity, depending on the degree of dissociation
  • Will produce solvonium ions, such as hydronium (H3O+) ions in water
  • Will denature proteins.[3]

pH scale

3-d model
Main Article: pH

You may have heard of something called the pH scale. This is what determines the hydrogen-ion acidity, of a solution; it is the negative logarithm of the hydrogen-ion concentration of a solution. In other words, any pH number higher than 7 is basic, and anything lower that 7 is acidic. Most PH values are whole numbers. [4]

Dissociation constant

An acid dissociation constant, Ka, (also known as acidity constant, or acid-ionization constant) is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction known as dissociation in the context of acid-base reactions. The equilibrium can be written symbolically as:

HA is in equilibrium with A + H+,

where HA is a generic acid that dissociates by splitting into A, known as the conjugate base of the acid, and the hydrogen ion or proton, H+, which, in the case of aqueous solutions, exists as a solvated hydronium ion. In the example shown in the figure, HA represents acetic acid, and A the acetate ion. The chemical species HA, A and H+ are said to be in equilibrium when their concentrations do not change with the passing of time. The dissociation constant is usually written as a quotient of the equilibrium concentrations (in mol/L), denoted by [HA], [A] and [H+]:

K_a equals the equilibrium concentration of the deprotonated form A-, times the equilibrium concentration of H+, all divided by the equilibrium concentration of the acid AH.

Due to the many orders of magnitude spanned by Ka values, a logarithmic measure of the acid dissociation constant is more commonly used in practice. pKa, which is equal to −log10 Ka, may also be referred to as an acid dissociation constant:

\ pK_{{{\mathrm  a}}}=-\log _{{10}}K_{{{\mathrm  a}}}