Acids, bases and salts

Introduction

Acids and bases play a central role in chemistry because, with the exception of redox reactions, every chemical reaction can be classified as an acid-base reaction. Our understanding of chemical reactions as acid-base interactions comes from the wide acceptance of the Lewis definition of acids and bases, which supplanted both the earlier Bronsted-Lowry concept and the first definition--the Arrhenius model. Arrhenius first defined acids as proton (H⁺) producers in aqueous solution and bases as hydroxide (OH^-) producers. Although this model is intuitively correct, it is limited to substances that include proton and hydroxide groups. Bronsted and Lowry proposed the more general definitions of acids and bases as proton donors and acceptors, respectively. Unlike the Arrhenius conception, the Bronsted-Lowry model accounts for acids in solvents other than water, where the proton transfers do not necessarily involve hydroxide ions. But the Bronsted-Lowry model fails to explain the observation that metal ions make water more acidic (discussed in Calculating pH's). Finally, Lewis gave us the more general definition of acids and bases that we use today. According to Lewis, acids are electron pair acceptors and bases are electron pair donors. Any chemical reaction that can be represented as a simple exchange of valence electron pairs to break and form bonds is therefore an acid-base reaction.

 Terms-Introduction

Acid  -  A substance that has the potential to donate a proton (H⁺) or accept an electron pair.

Acidic  -  Having a pH less than 7.

Arrhenius Model  -  Arrhenius proposed that acids are substances that produce protons, H⁺, in aqueous solution, whereas bases produce hydroxide ions, OH^-, in aqueous solution. Compare his model with the Bronsted-Lowry definition and the Lewis definition.

Base  -  A substance that can accept a proton, release OH^-, or donate an electron pair.

Bronsted-Lowry Definition  -  Bronsted and Lowry define an acid as a proton (H⁺) donor and a base as a proton acceptor. Compare this model with the Arrhenius Model and the Lewis definition.

Buffer  -  A solution composed of an acid and its conjugate base that serves to moderate the pH of the solution.

Conjugate Acid  -  A molecule that can be described as a base that has gained one proton.

Conjugate Base  -  A molecule that can be described as an acid that has lost one proton.

Indicator  -  A molecule whose conjugate acid or conjugate base has a different color. Indicators are used to measure the pH of a solution.

Lewis Definition  -  Lewis defined an acid as an electron pair acceptor and a base as an electron pair donor. Compare his model with the Arrhenius model and the Bronsted-Lowry definition.

pH  -  A measure of the hydrogen ion concentration, it is equal to - log [H⁺], where [H⁺] is the concentration of protons.

Redox  -  Short for "reduction-oxidation," a reaction that involves paired oxidation and reduction processes described in the Electrochemistry SparkNote.

Titration  -  An experiment that neutralizes an unknown amount of acid or base with a known volume and concentration of acid or base to determine the amount of unknown acid or base. 

 Fundamentals of acids and bases

 To achieve our goal of understanding how acids and bases work, we must first define what acids and bases are. There are three distinct conceptions of acids and bases that will be considered--the Arrhenius model, the Bronsted-Lowry model, and the Lewis model. Each of these models describe an acid-base reaction as a process by which a transfer occurs between two partner reagents, the acid and the base. For our purposes, the most useful model is the Bronsted-Lowry model, because we will be considering reactions involving proton transfers. Bronsted and Lowry described acids as proton donors and bases as proton acceptors. Calculations and measurements of pH are relevant to the Bronsted-Lowry conception of acid-base reactions. We will discuss each of the three models of acid-base reactions using representative equations. To speak of acid or base strength, we need scales for acidity and basicity. pH and pOH scales are quantitative representations of these values for acidic and basic solutions. Next, we will define the acid dissociation constant, K _a, and the base dissociation constant, K _b, to quantify the strengths of particular acids and bases. These terms allow us to process acid and base strength mathematically and package them into values that we can gage conceptually. Using reference values, we can then see that a solution with pH 6 is weakly acidic, since it is slightly lower in pH than water at pH 7.

 Terms- Fundamentals of acids and bases

Acid  -  A substance that has the potential to donate a proton or accept an electron pair.

Acidic  -  Having a pH less than 7 or a pOH greater than 7.

Amphiprotic  -  A species that can either donate or accept a proton, e. g. water.

Amphoteric  -  A species that can either donate or accept a hydroxide ion, such as Al(OH)₃. Many chemistry texts incorrectly use this term to mean that a substance can act as either an acid or a base.

Arrhenius Model  -  Arrhenius proposed that acids are substances that produce protons (H⁺) in aqueous solution, whereas bases produce hydroxide ions (OH^-) in aqueous solution. Compare his model with the Bronsted-Lowry definition and the Lewis definition.

Base  -  A substance that can accept a proton, release OH^-, or donate an electron pair.

Basic  -  Having a pH greater than 7 or a pOH less than 7.

Bronsted-Lowry Definition  -  Bronsted and Lowry define an acid as a proton (H⁺) donor and a base as a proton acceptor. Compare this model with the Arrhenius Model and the Lewis definition.

Conjugate Acid  -  A molecule that can be described as a base that has gained one proton.

Conjugate Base  -  A molecule that can be described as an acid that has lost one proton.

Dissociate  -  Separate into its ion constituents.

Lewis Definition  -  Lewis defined an acid as an electron pair acceptor and a base as an electron pair donor. Compare his model with the Arrhenius model and the Bronsted-Lowry definition.

pH  -  A measure of the hydrogen ion concentration, it is equal to - log [H⁺], where [H⁺] is the concentration of protons.

pK _a  -  A measure of the strength of an acid, it is equal to – log K _a, where K _a is the acid dissociation constant in water.

pK _b  -  A measure of the strength of a base, it is equal to – log K _b, where K _b is the base dissociation constant in water.

pOH  -  A measure of the hydroxide ion concentration, it is equal to - log [OH^-], where [OH^-] is the concentration of hydroxide ions.

Strong Acid  -  An acid with a pK _a less than zero. Strong acids completely dissociate in water.

Strong Base  -  A base with a pK _b less than zero. Strong bases completely dissociate in water.

Weak Acid  -  An acid with a pK _a greater than zero. Weak acids do not completely dissociate in water.

Weak Base  -  A base with a pK _b greater than zero. Weak bases do not completely dissociate in water.

Neutralization reaction 

 In chemistry, neutralization (or neutralisation, see spelling differences) is a chemical reaction in which an acid and a base react to form a salt. Water is frequently, but not necessarily, produced as well. Neutralizations with Arrhenius acids and bases always produce water where acid–alkali reactions produce water and a metal salt.Often, neutralization reactions are exothermic (the enthalpy of neutralization). For example, the reaction of sodium hydroxide and hydrochloric acid.

A neutralization reaction is a type of double displacement reaction. Typically, the resulting solution produced by the reaction consists of a salt and water. The general formula for acid–base neutralization reactions can be written as

acid + base → salt + water

HA + BOH → BA + H₂O

where HA represents the Arrhenius acid, BOH represents the Arrhenius base, and BA is the salt produced. Notice how, typical of a double replacement reaction, the cations and anions of the substances merely switch places.
An example reaction of this form is the reaction between hydrochloric acid (HCl) and sodium hydroxide (NaOH):

NaOH + HCl → NaCl + H₂O

Water and sodium chloride, or common table salt are produced.
The following are other examples of acid-base neutralisation reactions

• Sulphuric acid reacting with ammonium hydroxide to produce ammonium sulfate and water

H₂SO₄ + 2NH₄OH → (NH₄)₂SO₄ + 2H₂O

• Carbonic acid reacting with sodium hydroxide to produce sodium carbonate and water

H₂CO₃ + 2NaOH → Na₂CO₃ + 2H₂O

• Hydrochloric acid reacting with aluminium hydroxide to produce aluminium chloride and water

3HCl + Al(OH)₃ → AlCl₃ + 3H₂O

An Easy Trick To Find Products Of An Neutralisation Reaction 

To know this trick, you should understand one simple thing

H⁺ + OH^- = H₂O

For example,

Take HCl and NaOH

HCl + NaOH

a) Strike off the H in acid and OH in base

 HCl + NaOH

 

b) Add the remaining 

      

Cl + Na = NaCl

c) So the  products will be NaCl and H₂O

d) The reaction is HCl + NaOH = NaCl +  H₂O