Why acids produce hydrogen ions?

Question

where it's electron goes

Answer 1

One of the most popular Definition of an acid is "a compound that releases Proton or accept the electron" Proton is nothing but a Hydrogen atom without Electron. Acids do not give Hydrogen in all cases. When Acid reacts with Metal - the remaining part of Acid (for example Cl of Hydrochloric acid or SO4 of Sulphuric acid) combines with the metal and hydrogen is produced.
Fe + 2 HCl ---------> FeCl2 + 2 {H]
In no reaction Proton is released. It is always Hydrogen. Proton gets an electron from the metal in this case to form Hydrogen and accept the Cl- legend, as Cl- has one electron extra.
When Acid reacts with Alkali it forms Salt and Water.
HCl + NaOH --------> NaCl + H2O.

Answer 2

The electron goes onto the rest of the acid, making it a negative ion.

For example, HCl forms H+ and Cl-.

The acid produces H+ ions because the bond to the H is already weak.

Answer 3

The acid hydrochloric acid is monoprotic( one hydrogen)
When hydrogen chloride gas is dissolved in water, the hydrogen and the chloride components DISSOCIATE.
However, as chlorine is strongly electronegative it attracts an electron towards itself to become the chloride ion.(Cl-).
This electron that chlorine gains is lost from the hydrogen atom to become the hydrogen ion (H+).
All acids be they mineral or carboxylic dissociate to to some degree and the hydrogen's electron is lost to the conjugate anion.

Answer 4

An acid is by definition a compound that releases hydrogen ions when placed in solution.

Answer 5

All Acids have Hydrogen .
Hydrogen Behaves both as a ametal and non-metal
When we add a metal to an acid the hydrogen is displaced from the acid.
when electricity is passed the acid is ionized
fo ex:H2So4 --> H+ : So4 2-
HCl --> H+ : Cl-

Answer 6

acids generally dissoiates in aqous solutions like water to give hydrogen and respective ions because H+ions has more affinity to react with water to produce hydronium ions(H30).for example-
HCL + H2O =H3O+CL

Answer 7

Let us try to understand the Acids :
ARRHENIUS DEFINITION (1890)
This considered that acids were substances that released H+ ions in solution and that bases were substances that produced OH- ions in solution. It explained adequately how acids and bases neutralized each other

i.e. H+(aq) + OH‑(aq) -------> H2O(l)

acid base neutral

However, as time passed the Arrhenius definition began to seem less satisfactory. For one thing, there seemed to be two kinds of base. Metal hydroxides such as NaOH produce OH- ions in water by ionic dissociation. Bases, such as ammonia produce OH- ions in aqueous solution by undergoing a reaction with water.

NaOH(s) --------> Na+(aq) + OH-(aq)

NH3(g) + H2O(l) ------> NH+4(aq) + OH-(aq)



But more important, the Arrhenius definition of acids and bases was narrow.
It defined acids and bases in aqueous solution only.

THE LOWRY-BRONSTED DEFINITION (1923)

Lowry-Bronsted produced more general definitions of acids and bases.

These were in terms of H+ ions (called a protons, since a H+ ion has neither electrons nor neutrons).

According to Bronsted - an acid is a substance that can donate a proton

a base is a substance that can accept a proton


This definition can be represented by the general chemical reaction



A D B + H+



which does not attempt to show electrical charge balance.

In this equation -


· A is the acid.,

· B is the base and

· H+ (a hydrogen atom without an electron) is a proton.



Together A and B are called a CONJUGATE ACID AND BASE PAIR.

We can call B the conjugate base of A and call A the conjugate acid of B. This Lowry-Bronsted definition is more general than the Arrhenius definition. It does not refer to any specific solvent. We can simply say that any substance that can lose a proton is an acid.


Similarly the Bronsted bases are substances which can accept protons .





Mg(OH)+ + H+ = Mg2+ + H2O

While the relationship A----> B + H+ is a very good general definition of an acid and a base, there is a problem in applying this definition to an acid/base system in solution. The problem arises because the free protons (H+) cannot exist in solution to any great extent. In many cases the protons interact with the solvent



e.g.H2O(l) + H+(aq) ------> H3O+(aq) (a hydroxonium ion)


The solvent is then acting as a Bronsted base in accepting a proton.

The same reasoning can be applied to a base in solution. The base must obtain a proton from a proton source. Very often the proton source is the solvent.





The solvent is then acting as a Bronsted acid by providing a proton.



We can see that acid base reactions in solution are not simply processes in which an acid loses a proton or a base gains a proton. We must regard all acid/base reactions in solution as PROTON TRANSFER REACTIONS. Any such reaction includes TWO ACID - BASE CONJUGATE PAIRS. The original acid base pair, plus another conjugate pair, to accept the proton from the acid or donate the proton to the base. As we have said, this second acid-base pair is often derived from the solvent.


Each substance in the acid column is converted to its partner in the base column by the removal of a proton. Each substance in the base column is converted to its partner in the acid column by the addition of a proton.

ANY TWO CONJUGATE ACID/BASE PAIRS CAN BE C0MBINED IN AN ACID/BASE REACTION

e.g. HCl + OH- D H2O + Cl-



HCl and Cl- are one conjugate acid-base pair and H2O and OH- are another pair. The proton donated by the acid HCl is accepted by the base OH- in the forward reaction and a proton accepted by the Cl- is donated by the H2O in the reverse reaction.

The Bronsted definition can be summarized in a few sentences.

An acid is -

· a proton donor

· a base a proton acceptor.

We emphasize the relationship by designating conjugate acid base pairs.

e.g. HCl/Cl-
Acid base reactions in solution require two conjugate acid/base pairs because free protons (H+) do not exist in solution

THE LEWIS DEFINITION (1938)
The Lowry-Bronsted theory has its short-comings. Since it defines an acid as a proton donor, it excludes from the category of acids, substances that have no protons to donate. This limitation was overcome by a more general definition of acids and bases, BASED ON ELECTRONIC STRUCTURE, which was proposed by G N Lewis.

By the Lowry-Bronsted definition, a substance must accept a proton to be classified as a base. In other words the base must form a chemical bond with the proton (H+ ion). Since the proton has no electrons the base must have an electron pair available to form a bond.

In the Lewis definition:-

A BASE IS A SUBSTANCE THAT HAS A NON-BONDING VALENCE ELECTRON PAIR THAT CAN BE USED TO FORM A CHEMICAL BOND.

more simply - a Lewis base is an electron - pair donor.

The Lewis definition does not greatly expand our ideas about bases, but it does significantly broaden the category of substances that can be classified as acids. When a base accepts a proton it donates electrons to form the bond with the proton. Therefore the proton accepts the electrons.
A LEWIS ACID IS A SUBSTANCE THAT IS AN ELECTRON-PAIR ACCEPTOR

A proton (H+) is the simplest Lewis acid but many other substances fit the definition, including a number of cations, e.g Zn2+

Zn2+ + 2OH- ---> Zn(OH)2



Here the Zn2+ is the Lewis acid and OH- the Lewis base.

In organic chemistry it is especially common to find cations that behave as Lewis acids., accepting electrons. The reaction :-



C2H4 + Br2 ---> C2H4Br2

can be considered to proceed in 3 steps.
1. Br2 forms Br+ and Br ions. The Br+ is the Lewis acid.

2. This joins to C2H4 (a Lewis base).

3. In the third step Br- , a Lewis base, donates an electron pair and joins to the cation formed in the first step. The cation is a Lewis acid


H H H H

Br C C + + Br Br C C Br

H H H H



A great many reactions can be understood as the combination of a Lewis acid and a Lewis base. In fact a common method of classifying chemical reagents uses the Lewis acid and Lewis base definition.
A LEWIS ACID, WHICH IS A SUBSTANCE SEEKING ELECTRONS, IS CALLED AN ELECTROPHILE.

A LEWIS BASE, WHICH IS AN ELECTRON DONOR (OR A SUBSTANCE THAT SEEKS SUBSTANCES THAT ARE ELECTRON DEFICIENT), CAN BE CALLED A NUCLEOPHILE.

In the above reaction between C2H4 and Br2, Br+ is an electrophile and Br- is a nucleophile.

Many Lewis acids are substances that are chemically neutral. These substances can accept electrons and form additional bonds, either because they have incomplete octets or octet expansion is possible.
Many compounds of the elements of Group III are Lewis acids, the best known being the halides such as BF3, BCl3, AlCl3

Lewis acid (LA) can accept a pair of electrons and form a coordinate covalent bond. The Lewis acid and Lewis base theory, named after the American chemist Gilbert Lewis, is one of several acid-base reaction theories. Therefore the term acid, per se, is ambiguous; it should always be clarified as being a Lewis acid or a Brønsted-Lowry acid.

An electrophile or electron acceptor is a Lewis acid. A Lewis acid usually has a low-energy LUMO, which interacts with the HOMO of the Lewis base. Unlike a Brønsted-Lowry acid, which always transfers a hydrogen ion (H+), a Lewis acid can be any electrophile (including H+). Although all Brønsted-Lowry acids are Lewis acids, in common usage the term Lewis acid is often reserved for those Lewis acids which are not Brønsted-Lowry acids.

The reactivity of Lewis acids can be judged from the Hard-Soft Acid-Base concept. There is no universally valid description of Lewis acid strength, because Lewis acid strength depends on the specific Lewis base. Christe and Dixon[1] have predicted Lewis acid strength based on a computational model of gas-phase affinity for fluoride, and out of a selection of common isolable Lewis acids they found that SbF5 had the strongest fluoride affinity. Fluoride is a "hard" Lewis base; chloride and "softer" Lewis bases are very difficult to study because of limitations of the computational methods, and Lewis acidity in solution suffers from the same restriction.[2]

Some common Lewis acids include aluminium chloride, iron(III) chloride, boron trifluoride, niobium pentachloride and the lanthanide triflates such as ytterbium(III) triflate.

Answer 8

The electron goes onto the rest of the acid, making it a negative ion.

For example, HCl forms H+ and Cl-.

The acid produces H+ ions because the bond to the H is already weak.


An acid is by definition a compound that releases hydrogen ions when placed in solution.

The acid hydrochloric acid is monoprotic( one hydrogen)
When hydrogen chloride gas is dissolved in water, the hydrogen and the chloride components DISSOCIATE.
However, as chlorine is strongly electronegative it attracts an electron towards itself to become the chloride ion.(Cl-).
This electron that chlorine gains is lost from the hydrogen atom to become the hydrogen ion (H+).
All acids be they mineral or carboxylic dissociate to to some degree and the hydrogen's electron is lost to the conjugate anion.

One of the most popular Definition of an acid is "a compound that releases Proton or accept the electron" Proton is nothing but a Hydrogen atom without Electron. Acids do not give Hydrogen in all cases. When Acid reacts with Metal - the remaining part of Acid (for example Cl of Hydrochloric acid or SO4 of Sulphuric acid) combines with the metal and hydrogen is produced.
Fe + 2 HCl ---------> FeCl2 + 2 {H]
In no reaction Proton is released. It is always Hydrogen. Proton gets an electron from the metal in this case to form Hydrogen and accept the Cl- legend, as Cl- has one electron extra.
When Acid reacts with Alkali it forms Salt and Water.
HCl + NaOH --------> NaCl + H2O.

Let us try to understand the Acids :
ARRHENIUS DEFINITION (1890)
This considered that acids were substances that released H+ ions in solution and that bases were substances that produced OH- ions in solution. It explained adequately how acids and bases neutralized each other

i.e. H+(aq) + OH‑(aq) -------> H2O(l)

acid base neutral

However, as time passed the Arrhenius definition began to seem less satisfactory. For one thing, there seemed to be two kinds of base. Metal hydroxides such as NaOH produce OH- ions in water by ionic dissociation. Bases, such as ammonia produce OH- ions in aqueous solution by undergoing a reaction with water.

NaOH(s) --------> Na+(aq) + OH-(aq)

NH3(g) + H2O(l) ------> NH+4(aq) + OH-(aq)



But more important, the Arrhenius definition of acids and bases was narrow.
It defined acids and bases in aqueous solution only.

THE LOWRY-BRONSTED DEFINITION (1923)

Lowry-Bronsted produced more general definitions of acids and bases.

These were in terms of H+ ions (called a protons, since a H+ ion has neither electrons nor neutrons).

According to Bronsted - an acid is a substance that can donate a proton

a base is a substance that can accept a proton


This definition can be represented by the general chemical reaction



A D B + H+



which does not attempt to show electrical charge balance.

In this equation -


· A is the acid.,

· B is the base and

· H+ (a hydrogen atom without an electron) is a proton.



Together A and B are called a CONJUGATE ACID AND BASE PAIR.

We can call B the conjugate base of A and call A the conjugate acid of B. This Lowry-Bronsted definition is more general than the Arrhenius definition. It does not refer to any specific solvent. We can simply say that any substance that can lose a proton is an acid.


Similarly the Bronsted bases are substances which can accept protons .





Mg(OH)+ + H+ = Mg2+ + H2O

While the relationship A----> B + H+ is a very good general definition of an acid and a base, there is a problem in applying this definition to an acid/base system in solution. The problem arises because the free protons (H+) cannot exist in solution to any great extent. In many cases the protons interact with the solvent



e.g.H2O(l) + H+(aq) ------> H3O+(aq) (a hydroxonium ion)


The solvent is then acting as a Bronsted base in accepting a proton.

The same reasoning can be applied to a base in solution. The base must obtain a proton from a proton source. Very often the proton source is the solvent.





The solvent is then acting as a Bronsted acid by providing a proton.



We can see that acid base reactions in solution are not simply processes in which an acid loses a proton or a base gains a proton. We must regard all acid/base reactions in solution as PROTON TRANSFER REACTIONS. Any such reaction includes TWO ACID - BASE CONJUGATE PAIRS. The original acid base pair, plus another conjugate pair, to accept the proton from the acid or donate the proton to the base. As we have said, this second acid-base pair is often derived from the solvent.


Each substance in the acid column is converted to its partner in the base column by the removal of a proton. Each substance in the base column is converted to its partner in the acid column by the addition of a proton.

ANY TWO CONJUGATE ACID/BASE PAIRS CAN BE C0MBINED IN AN ACID/BASE REACTION

e.g. HCl + OH- D H2O + Cl-



HCl and Cl- are one conjugate acid-base pair and H2O and OH- are another pair. The proton donated by the acid HCl is accepted by the base OH- in the forward reaction and a proton accepted by the Cl- is donated by the H2O in the reverse reaction.

The Bronsted definition can be summarized in a few sentences.

An acid is -

· a proton donor

· a base a proton acceptor.

We emphasize the relationship by designating conjugate acid base pairs.

e.g. HCl/Cl-
Acid base reactions in solution require two conjugate acid/base pairs because free protons (H+) do not exist in solution

THE LEWIS DEFINITION (1938)
The Lowry-Bronsted theory has its short-comings. Since it defines an acid as a proton donor, it excludes from the category of acids, substances that have no protons to donate. This limitation was overcome by a more general definition of acids and bases, BASED ON ELECTRONIC STRUCTURE, which was proposed by G N Lewis.

By the Lowry-Bronsted definition, a substance must accept a proton to be classified as a base. In other words the base must form a chemical bond with the proton (H+ ion). Since the proton has no electrons the base must have an electron pair available to form a bond.

In the Lewis definition:-

A BASE IS A SUBSTANCE THAT HAS A NON-BONDING VALENCE ELECTRON PAIR THAT CAN BE USED TO FORM A CHEMICAL BOND.

more simply - a Lewis base is an electron - pair donor.

The Lewis definition does not greatly expand our ideas about bases, but it does significantly broaden the category of substances that can be classified as acids. When a base accepts a proton it donates electrons to form the bond with the proton. Therefore the proton accepts the electrons.
A LEWIS ACID IS A SUBSTANCE THAT IS AN ELECTRON-PAIR ACCEPTOR

A proton (H+) is the simplest Lewis acid but many other substances fit the definition, including a number of cations, e.g Zn2+

Zn2+ + 2OH- ---> Zn(OH)2



Here the Zn2+ is the Lewis acid and OH- the Lewis base.

In organic chemistry it is especially common to find cations that behave as Lewis acids., accepting electrons. The reaction :-



C2H4 + Br2 ---> C2H4Br2

can be considered to proceed in 3 steps.
1. Br2 forms Br+ and Br ions. The Br+ is the Lewis acid.

2. This joins to C2H4 (a Lewis base).

3. In the third step Br- , a Lewis base, donates an electron pair and joins to the cation formed in the first step. The cation is a Lewis acid


H H H H

Br C C + + Br Br C C Br

H H H H



A great many reactions can be understood as the combination of a Lewis acid and a Lewis base. In fact a common method of classifying chemical reagents uses the Lewis acid and Lewis base definition.
A LEWIS ACID, WHICH IS A SUBSTANCE SEEKING ELECTRONS, IS CALLED AN ELECTROPHILE.

A LEWIS BASE, WHICH IS AN ELECTRON DONOR (OR A SUBSTANCE THAT SEEKS SUBSTANCES THAT ARE ELECTRON DEFICIENT), CAN BE CALLED A NUCLEOPHILE.

In the above reaction between C2H4 and Br2, Br+ is an electrophile and Br- is a nucleophile.

Many Lewis acids are substances that are chemically neutral. These substances can accept electrons and form additional bonds, either because they have incomplete octets or octet expansion is possible.
Many compounds of the elements of Group III are Lewis acids, the best known being the halides such as BF3, BCl3, AlCl3

Lewis acid (LA) can accept a pair of electrons and form a coordinate covalent bond. The Lewis acid and Lewis base theory, named after the American chemist Gilbert Lewis, is one of several acid-base reaction theories. Therefore the term acid, per se, is ambiguous; it should always be clarified as being a Lewis acid or a Brønsted-Lowry acid.

An electrophile or electron acceptor is a Lewis acid. A Lewis acid usually has a low-energy LUMO, which interacts with the HOMO of the Lewis base. Unlike a Brønsted-Lowry acid, which always transfers a hydrogen ion (H+), a Lewis acid can be any electrophile (including H+). Although all Brønsted-Lowry acids are Lewis acids, in common usage the term Lewis acid is often reserved for those Lewis acids which are not Brønsted-Lowry acids.

The reactivity of Lewis acids can be judged from the Hard-Soft Acid-Base concept. There is no universally valid description of Lewis acid strength, because Lewis acid strength depends on the specific Lewis base. Christe and Dixon[1] have predicted Lewis acid strength based on a computational model of gas-phase affinity for fluoride, and out of a selection of common isolable Lewis acids they found that SbF5 had the strongest fluoride affinity. Fluoride is a "hard" Lewis base; chloride and "softer" Lewis bases are very difficult to study because of limitations of the computational methods, and Lewis acidity in solution suffers from the same restriction.[2]

Some common Lewis acids include aluminium chloride, iron(III) chloride, boron trifluoride, niobium pentachloride and the lanthanide triflates such as ytterbium(III) triflate
All Acids have Hydrogen .
Hydrogen Behaves both as a ametal and non-metal
When we add a metal to an acid the hydrogen is displaced from the acid.
when electricity is passed the acid is ionized
fo ex:H2So4 --> H+ : So4 2-
HCl --> H+ : Cl-