Acid base and salt in chemistry

Acid base and salt in chemistry

This post would cover the basics of acid base and salt in chemistry.  This post would cover acid and base reactions such as acid base neutralization reaction, reactions of acid with metal etc. Also, important topics such as acid and base definition, arrhenius acid and base, bronsted-Lowry theory and measuring acid base ph will be covered.

Acids

Now let us get the concepts of acids and bases crystal clear

The water molecules may break down to H⁺ and OH^– ions as shown in the following equation below:

H₂O ⇆ H⁺ + OH^–

àThe H⁺ and OH^– also combine to form H₂O molecule, the thing is under equilibrium.

• When water molecule breaks down it releases an equal number of H⁺ and OH^– ions, when the ions combine, both of them reduce equally.
• The main point is that without any outside interference the concentration of the H⁺ and OH^– ions should be equal.
• In a neutral solution, the concentration of H⁺ and OH^– ions are equal.
• On the other hand in an acidic solution, the concentration of H⁺ ions is more than the OH^– ions, more H⁺ ions than OH^– The higher the concentration of H⁺ than the OH-, the more acidic the solution is.
• In an alkaline solution, the concentration of H⁺ ions is less than the OH^– More OH^– ions than the H⁺ ions. The higher the concentration of OH^– than the H⁺ ions, the more alkaline the solution is.

To measure the acidity or alkaline values we use a scale, called the pH scale. The pH scale is between 0 to 14.

pH value = 7 , means the solution is neutral

If pH value less than 7 means the solution is acidic, the lower the pH value than 7, the more acidic the solution is.

pH value more than 7 means the solution is alkaline, the higher the pH value than 7, the more alkaline the solution is.

To get an even better understanding of acid-base concepts you can check out this book “Acids and Bases (Why Chemistry Matters)” on Amazon.

Measuring pH Value:

There are three ways to measure the pH values:

By using universal indicator

the universal indicator gives various colors for different pH values. It can be in the form of paper or solution. But it’s not that accurate. A chart with marked color is used to find out the pH values.

By using a pH meter

It’s a device which simply gives the pH value when its probe is dipped in a solution.

Simple Indicators

These kinds of indicators give only three different colors in acidic, neutral and alkaline solutions. These are to indicate the end point of a reaction when an acid or alkali is neutralized and a color change is seen, these indicators give sharp color change. Universal indicators are not good for indicating the end point because it does not give a sharp color change. There are usually three common simple indicators:

The reaction of the acid with metals:

Metals below hydrogen in the reactivity series do not react with acid. Eg: copper, silver, gold etc

Metal above hydrogen in the reactivity series reacts with acid to form a salt and hydrogen gas.

Metal + Acid –> Salt + Hydrogen

Mg + 2HCl –> MgCl₂ + H₂

• Highly reactive metals react too much violently with acid, so we usually do not react such metals with acid. E.g: sodium, potassium, lithium etc.
• To get the desired salt, we must react the acid with the following acids.

Reaction of magnesium with acids

Observation: when magnesium reacts with the acid, a rapid reaction occurs. Fizzing occurs, colorless gas evolved which relights a glowing splint with a pop sound (hydrogen)

• Reaction with sulfuric acid: Mg(s) + H₂SO₄(aq) –> MgSO₄(aq) + H₂(g)
• Reactions with hydrochloric acid: Mg(s) + HCl(aq) –> MgCl₂ (aq) + H₂(g)
• Type of reaction of magnesium with the acids: displacement reaction, the more reactive magnesium displaces the replaces the less reactive hydrogen.
• Why do the reactions look so similar?

Answer: it becomes clear when we write the ionic equation of the above reactions with acid

For Sulfuric acid:

Mg(s) + 2H⁺(aq) + SO₄^2-(aq) –> Mg²⁺(aq) + SO₄^2-(aq) + H₂(g)

Ionic equation (by omitting the spectator ion SO₄^2-) :  Mg(s) + 2H⁺(aq) –> Mg²⁺(aq) + H₂(g)

For Hydrochloric acid:

Mg(s) + 2H⁺(aq) + 2Cl^–(aq) –> Mg²⁺(aq) + 2Cl^–(aq) + H₂(g)

Ionic equation (by omitting the spectator ion Cl^–) : Mg(s) + 2H⁺(aq) –> Mg²⁺(aq) + H₂(g)

• By looking at the reactions, we can clearly see that they are same. The magnesium reacts with the H⁺
• Reaction of zinc with acids: zinc also reacts with acids in a similar fashion as magnesium, but the reaction is slower. Because zinc is less reactive than magnesium. The reaction can be speeded up, by heating the acid and zinc mixture, or by adding copper(II)sulfate solution to it.

Zn(s) + H₂SO₄(aq) –> ZnSO₄(aq) + H₂(g)

Zn(s)+ 2HCl(aq) –> ZnCl₂(aq) + H₂(g)

Ionic equation: Zn(s)+ 2H⁺(aq) –> Zn²⁺(aq) + H₂(g)

Making Hydrogen gas in lab

To make hydrogen gas in the lab, the zinc metal is usually reacted with dilute sulfuric acid, with little amount of copper (II)sulfate solution added to speed up reaction.

• Test for hydrogen: relights a glowing splint with a pop sound.

Reaction that occurs: 2H₂(g) + O_{2 –>} 2H₂O(l)

Reacting Acids with metal oxides

Acid + Metal Oxide à Salt + Water

Example:

Copper (II) Oxide + Sulfuric acid à Salt + water

CuO(s) + H₂SO₄ (aq) –> CuSO₄ (aq) + H₂O (l)

Observation: Black copper (II) Oxide reacts with colorless sulfuric acid to produce a blue solution of copper (II) Sulfate.

• Writing the ionic equation for the above reaction:

1. Firstly breakdown the above equations into ions:

Cu²⁺(s) + O^2-(s) + 2H⁺(aq) + SO₄^2-(aq) –>  Cu²⁺(aq) + SO₄^2-(aq) + H₂O (l)

1. Here the Cu²⁺ and SO₄^2- are spectator ions as they are unchanged on both sides:

Writing the final ionic equation without the spectator ions:

O^2-(s) + 2H⁺(aq) –> H₂O(l)

• Bases: bases are defined as compounds which combine with hydrogen ions.

Reaction of acid with metal hydroxide

Acid + metal hydroxide à Salt + water

Example:

Hydrochloric acid + sodium hydroxide à sodium chloride + water

HCl(aq) + NaOH (aq) –> NaCl (aq) + H₂O (l)

Writing ionic equation:

1. H⁺(aq) + Cl^–(aq) + Na⁺(aq) + OH^–(aq) –> Na⁺(aq) + Cl^–(aq) + H₂O (l)
2. Writing the final ionic equation without spectator ions Na⁺ and Cl^–

H⁺(aq) + OH^–(aq) –> H₂O (l)

Reaction of acid with carbonates (metal carbonates)

Carbonate + acid –> salt + water + carbon dioxide gas

Example:

Copper(II) carbonate + Sulfuric acid à  Copper(II)Sulfate + water + carbon dioxide

CuCO₃(s)+ H₂SO₄(aq) –> CuSO₄(aq) + H₂O (l) + CO₂ (g)

Writing ionic equation:

1. Cu²⁺(s) + CO₃^2- (s) + 2H⁺(aq) + SO₄^2-(aq) –> Cu²⁺(aq) + SO₄^2-(aq) + H₂O (l) + CO₂ (g)
2. Writing the final ionic equation without the spectator ions Cu²⁺ and SO₄^2-

CO₃^2- (s) + 2H⁺(aq) –> H₂O (l) + CO₂ (g)

  More Examples:

CuCO₃(s) + 2HNO₃ (aq) –> Cu(NO₃)₂ (aq) + CO₂(g) + H₂O(l)

CuCO₃(s) + 2HCl(aq) –> CuCl₂ (aq) + CO₂ (g) + H₂O(l)

Arrhenius theory and Bronsted-Lowry theory

According to Arrhenius theory:

• An acid is a substance which releases H⁺ ions when dissolved in water.
• Bases are solutions containing OH^– ions

Arrhenius theory is very limited and fails to explain some situations. So another theory came out, Bronsted-Lowry theory.

According to Bronsted-Lowry theory:

• An acid is a proton donor
• A base is a proton acceptor

In this theory, the H⁺ ion is referred to as a proton. Because a hydrogen atom after losing the single electron in the outer shell is only a proton.

For example the reaction between ammonia and hydrogen chloride gas:

NH₃(g) + HCl(g) –> NH₄Cl

NH₃(g) + HCl(g) –> NH₄⁺ + Cl^–

Here the ammonia molecule gains a proton (H⁺ ion) hence it’s a base, while the HCl looses a proton (H⁺ ion) hence it’s an acid.

H⁺ ions do not exist in solutions singularly, but as the H₃O⁺ ion (hydroxonium ion), but we use the simplified H⁺ version only.

The example of an acid with water:

HCl (g) + H₂O(l) –> H₃O⁺(aq) + Cl^–

Here HCl acts as an acid as it donates a proton, while the H₂O acts as base as it gains the proton.