Learn All About The Strong Acids and Bases

Acids and Bases play a vital role in our life and all of the environment around us. When you brush your teeth with toothpaste or wash your hands with soap or take your coffee, you’ve already experienced acids and bases.

In chemistry, substances are classified as acids or bases or neutral, depending on some characteristics like PH and behavior with water. Acids and bases are also sub classified into weak and strong acids and bases, depending on the strength. In this article we will focus on strong acids and bases but first we will clarify some important definitions.

General Look at Acids and Bases

Acids and Bases

Before explaining strong acids and bases, we will first clarify what is meant by acids and bases in general and based on:

  •       Hydrogen ion H+ or hydroxide ion OH- production when dissolved in water.
  •        Transfer of proton.

In case of mixing acids and bases, they react together to neutralize the properties and forming a salt and water.

mixing acids and bases

Acids

An acid is a matter that produces ions of hydrogen (H+) when dissolved in water. All acids must contain at least one hydrogen ion. Acids are Protons donors.

Ex:-                             HCl (aq) ———– H+ (aq) + Cl (aq)

 Acids must have acidic protons (H+) to be donated and depending on no of acidic protons, acids can be

          Monoprotic (ex: hydrochloric acid and nitric acid)

          Diprotic (ex: sulfuric acid)

          Triprotic (ex: phosphoric acid)

Water can act as an acid when donates a proton to a base and forms hydroxide ion (OH).  

Note:

Super acid is acid with more acidity than pure sulfuric acid. Some are capable of protonating nearly anything, including hydrocarbons. An example of super acids is fluorosulfonic acid HSO3F.

Bases

A base is a matter that neutralizes acid. Bases split to form ions of hydroxide (OH) when dissolved in water ( the base that has been added to water is called an alkaline solution). Bases are protons acceptors and to accept protons, they must have a lone pair.

Ex:-                                                      NH3 + H+ —————————– NH4+

Bases can be neutral (ex: ammonia) or negatively charged (ex: hydroxide ion).

Water can act as a base when it accepts a proton from an acid and forms hydronium ion H3O+.

Examples of acids and bases (household):

house hold acids and bases

Properties of Acids and Bases

The following table shows the main properties of acids and bases:

Acids

Bases

Acids taste sour. Bases taste bitter.
React with some metals to separate hydrogen. Bases feel soupy to touch as they can Change the protein structure.
React with bases and form salts. React with acids and form salts.
Change the color of litmus paper from blue to red. Change the color of litmus paper from red to blue.
Good electrical current conductor. Good electrical current conductor.
Can destroy cells. Strong bases can destroy the skin proteins and cause severe burns.
Can catalyze chemical reactions. Corrosive substances that can damage other substances.
Soluble in water. Some bases are soluble in water but others are not.
Mineral acids are colorless and in solid case, organic acids are white. Most of bases are colorless.
Commonly used in car batteries and fertilizers. Commonly used in several cleaning products.
Examples: Vinegar (acetic acid) /  Hydrochloric acid / Sulfuric acid Examples: Ammonia / Sodium hydroxide / Potassium hydroxide

Difference between Acids and Bases

 The main difference between acids and bases is the PH value of each of them. The acids pH are below 7 and the bases pH are above 7.

 The pH or potential of hydrogen is a measurement for the acidity or alkalinity of a solution on a logarithmic scale on which 7 is neutral, higher values are more alkaline and lower values are more acidic. The pH is equal to −log10 c, where c refers to the hydrogen ion concentration in moles per litre.

 Acids and Bases Identification and Measurement

When we test a matter with a pH meter, we get a number range from zero to 14. This is a pH scale, and it can be used to compare substances, also identify and measure acids and bases. note that the pH scale is logarithmic which means that 1 decrease in the pH scale can cause 10 times increase of hydrogen ions concentration.

pH scale

Acids have a pH value below 7. The more H+ ions, the more the acidity of it and the lower the pH value will be. Bases have a pH value above 7. When there is a balance of H+ and OH ions, it is called neutral and pH value is 7.

For very strong acids, the pH value can be less than zero and for very strong bases, the pH value can be more than 14.

The PH Scale

pH scale

  • pH 0 – 2: strong acid
  • pH 3 – 6: weak acid
  • pH 7: neutral
  • pH 8 – 11: weak alkali
  • pH 12 – 14: strong alkali

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Acids and Bases Examples

The following table shows list of acids and bases

Acids

Bases

   H2SO4 –  Sulfuric acid    KOH – Potassium hydroxide
   HCl – Hydrochloric  acid    NaOH – Sodium hydroxide
   CH3COOH – Acetic acid    Li(OH)2 – Lithium hydroxide
   HNO3 – Nitric acid    Ca(OH)2 – Calcium hydroxide
   HBr – Hydrobromic acid    Ba(OH)2 – Barium hydroxide
  H2CO3 – Carbonic acid    C5H5N – Pyridine
   HI – Hydroiodic acid    HCO3 – Bicarbonate ion
   H2CO3 – Carbonic acid    NH3 – Ammonia

Strong Acids and Bases

After understanding the general acids and bases differences, now we will clarify what is meant by strong acids and bases and their general properties.

Strong acids or alkalis  are nearly or completely ionised in water. Examples of strong acids (hydrochloric acid and sulphuric acid) / examples of Strong alkalis (sodium hydroxide and potassium).

Strong Acids

Strong acid is characterized by their ability to be completely ionized or dissociated in the solvents. The pKa value of strong acids are less than -2.

Note: when strong acids become more concentrated, they may be unable to fully dissolve.  A strong acid is fully dissociated in solutions of 1.0 M or lower concentration.

The general form of the dissociation or ionization reaction of the strong acids is:

HA + S ↔ SH+ + A-

 Where:

S is the solvent molecule like water

For example, hydrochloric acid dissociation in water:

HCl(aq)    →   H+(aq) + Cl(aq)

Ka is the acid dissociation constant and it measures how completely an acid dissociates in a solution. The large value of Ka indicates that the acid is highly dissociated into its ions and it is a strong acid. Strong acids have a small value of Ka as they completely dissociate in water.

There is a direct relationship between Ka and pKa (the logarithmic acid dissociation constant). Therefore, the stronger the acid the lower the pKa value.

pKa = -log [Ka]

 The value of pKa of an acid depends on the solvent. For example, the pKa value of hydrochloric acid is about -5.9 in water and in dimethyl sulfoxide is about -2.0, while the  pKa value of hydrobromic acid is about -8.8 in water and in dimethyl sulfoxide is about -6.8.

 

Strong Bases

Strong base is characterized by their ability to be completely ionized or dissociated into their respective ions.

Most common strong bases are metal hydroxides as they are composed of a metal ionically bonding with a hydroxyl ion OH−, like potassium hydroxide KOH.

Strong bases are considered as good acceptors of protons that cannot remain in aqueous solution. For instance, all oxygen ions are converted to hydroxide ions by accepting protons from water H2O molecules. As a result, water molecules are converted to hydroxide OH.note that the strong bases cations are soluble in water (if a cation is soluble in water, it can form a strong base).

Ex: BOH + H2O → B+(aq) + OH(aq)

The potential of hydroxide ions or pOH is a measurement of the concentration of hydroxide ion (OH-). The higher the hydroxide ions concentration in the solution, the lower the value of  pOH therefore, a stronger base.

pOH = -log [OH]

Strong bases commonly have a pH value between 13 \ 14.

Kb or the base dissociation constant measures how completely a base  ionizes or dissociates in a solution. The large value of Kb indicates a strong base and it also means the base is highly dissociated into its ions.

There is a direct relationship between Kb and pKb (the logarithmic base dissociation constant). Therefore, the stronger the acid the lower the pKb value.

pKb = -log [Kb]

The following equation shows the relation between pH and pOH of an aqueous solution:

pH + pOH = 14

If either the pOH or the pH value of a solution is known, you can calculate the other.

Examples of Acids and Bases

Examples of strong and weak acids:

List of Strong Acids and Bases

There are examples of common strong acids and strong bases.

Strong Acids:

  • H2SO4 (sulfuric acid)
  • HNO3 (nitric acid)
  • HCl (hydrochloric acid)
  • HI (hydroiodic acid)
  • HBr (hydrobromic acid)
  • HClO3 (chloric acid)
  • HClO4 (perchloric acid)
  • Triflic acid (H[CF3SO3])
  •   Fluoroantimonic acid (H[SbF6]).

 Strong Bases:

  • KOH (potassium hydroxide)
  • LiOH (lithium hydroxide)
  • NaOH (sodium hydroxide)
  • Ca(OH)2 (calcium hydroxide)
  • Ba(OH)2 (barium hydroxide)
  • RbOH (rubidium hydroxide)
  • Sr(OH)2 (strontium hydroxide)
  • CsOH (cesium hydroxide)

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Strong Acids and Bases Chart

Acid Base
Strength of Acid Ka Name Formula Formula Name Strength of Base
Strongest Large Perchloric acid HClO4 ClO4 Perchlorate
ion
Weakest
3.2 * 109 Hydroiodic acid HI I- Iodide
1.0 * 109 Hydrobromic acid HBr Br- Bromide
1.3 * 106 Hydrochloric acid HCl Cl- Chloride
1.0 * 103 Sulfuric acid H2SO4 HSO4 Hydrogen
sulfate ion
2.4 * 101 Nitric acid HNO3 NO3 Nitrate ion
——– Hydronium ion H3O+ H2O Water
5.4 * 10-2 Oxalic acid HO2C2O2H HO2C2O2 Hydrogen
oxalate ion
1.3 * 10-2 Sulfurous acid H2SO3 HSO3 Hydrogen
sulfite ion
1.0 * 10-2 Hydrogen sulfate ion HSO4 – SO4 2- Sulfate ion
7.1 * 10-3 Phosphoric acid H3PO4 H2PO4 Dihydrogen
phosphate ion
7.2 * 10-4 Nitrous acid HNO2 NO2 Nitrite ion
6.6 * 10-4 Hydrofluoric acid HF F- Fluoride ion
1.8 * 10-4 Methanoic acid HCO2H HCO2 Methanoate
ion
6.3 * 10-5 Benzoic acid C6H5COOH C6H5COO- Benzoate ion
Medium 5.4 * 10-5 Hydrogen oxalate ion HO2C2O2 O2C2O2 2- Oxalate ion Medium
1.8 * 10-5 Ethanoic acid CH3COOH CH3COO Ethanoate
(acetate) ion
4.4 * 10-7 Carbonic acid H2CO3 HCO Hydrogen
carbonate ion
1.1 * 10-7 Hydrosulfuric acid H2S HSO3 Hydrogen
sulfide ion
6.3 * 10-8 Dihydrogen phosphate ion H2PO HPO4 2- Hydrogen
phosphate ion
6.2 * 10-8 Hydrogen sulfite ion HS- SO32- Sulfite ion
2.9 * 10-8 Hypochlorous acid HClO ClO Hypochlorite
ion
6.2 * 10-10 Hydrocyanic acid HCN CN Cyanide ion
5.8 * 10-10 Ammonium ion NH4 + NH3 Ammonia
5.8 * 10-10 Boric acid H3BO3 H2BO3 Dihydrogenborate
4.7 * 10-11 Hydrogen carbonate ion HCO3  CO3 2- Carbonate
ion
4.2 * 10-13 Hydrogen phosphate ion HPO4 2- PO4 3- Phosphate
ion
1.8 * 10-13 Dihydrogen borate ion H2BO3  HBO3 2- Hydrogen
borate ion
1.3 * 10-13 Hydrogen sulfide ion HS- 2- Sulfide ion
1.6 * 10-14 Hydrogen borate ion HBO3 2- BO3 3- Borate ion
Weakest Water H2O OH- Hydroxide Strongest

The previous table shows the strong acids and bases strength chart.

Source: Merck

Properties of Strong Acids and Bases

Strong Acids

The most common properties of strong acids are:

  • They can release H+ ions in aqueous solutions (Arrhenius acid).

Arrhenius acid

(Bronsted-Lowry acid

  • They are able to receive electrons from a base (Lewis acid).

(Lewis acid)

  • Can catalyze chemical reaction.
  • able to conduct electricity.
  • The hydronium ions concentration is equal to the concentration of the acid.
  • They are highly reactive and dangerous, so they must be handled carefully.

Strong Bases

The most common properties of strong bases :

  • They are good proton (hydrogen ion) acceptors and electron donors.
  • They can deprotonate weak acids.
  • Their aqueous solutions are slippery and soapy.
  • They are less reactive and violent than acids.
  • They undergo many chemical reactions, especially with organic compounds (ex: the most common reaction is saponification).

Difference between Strong Acids and Bases

strong acids and bases

The main difference between strong acids and bases is that after the complete dissolving, strong acids give more hydrogen ions and the strong bases give more hydroxide ions in the solutions.

Weak Acids and Bases

weak acids and bases are partially ionised in water. Examples of weak acids (citric acid and carbonic acid) / examples of weak alkalis (copper hydroxide and ammonia).

when  weak acids are dissolved in water, an equilibrium is happened between the weak acid concentrations and its constituent ions, so we can said that weak acids are not completely ionized in solution.

For example, hydrofluoric acid HF is considered as weak acid. When dissolved in water, its ions exist in equilibrium with hydrogen ions, which reacts with water to form hydronium, and  flour ions. because the acids don’t completely dissociate into their ionic components, they are called weak acid in this case.

HA + H2O H3O+ + A

The concept is the same in case of weak base; but, in weak base case, we can define a base that cannot completely ionize or accept hydrogen ions which found  in the solution.

The base needs to be dissolved in solution to react with the acid and form an acid base pair. note that the base should break into its  ions to react in the way it must.

When weak base dissolved in water, the solution so, in this case we said that the weak base does not fully ionize as when it dissociated in water, the result solution contains a small amount of OH ions and a big amount of the nondissociated base.

Ex: the counterpart’s of the base in ammonium hydroxide (NH4OH) are ion of ammonia  NH4+ and hydroxide ion OH; the ion of ammonia NH4+  accepts the acidic hydrogen in solution; but, in case of  weak base these ion do not dissolve completely so there is a mixture of the acid base pair and ammonium hydroxide NH4OH.

B + H2O BH+ + OH

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