January 2019

Oxidation number and concept

    The oxidation number of CH4, CH3Cl, CH2Cl2, CHCl3, CCl4 and CO2 are -4, -2, 0, +2, +4 and +4 respectively. In H2, the oxidation number of hydrogen is zero but in H2O it is +1. Similarly, magnesium in the elementary state has a zero oxidation number but in MgCl2, the oxidation number is +2.
    From the above, we can define Oxidation and Reduction according to the Oxidation number increase or decrees,

Oxidation

    Oxidation may be defined as a process in which an increase in oxidation number by atoms or ions occurs. A reagent which can increase the oxidation number of an element or ion is called an oxidizing agent. 

Reduction

    The reduction is a process in which a decrease in oxidation number by atoms or ions occurs. A reagent that lowers the oxidation number of an element or ion is called a reducing agent.

Representation of oxidation and reduction

    Oxidation and reduction are always found to go hand to hand during a redox reaction. Whenever an element or a compound is oxidized, another element or another compound must be simultaneously reduced.
    An oxidant is reduced and simultaneously the reductant is oxidized.
Oxidation reduction reaction and oxidation number
Oxidation number and oxidation-reduction
    H2S + Cl2 → 2HCl + S
  1. Magnesium metal burns in oxygen to produce magnesium oxide: Before the reaction oxidation number of magnesium and oxygen both zero. But after the reaction, the oxidation number of magnesium and oxygen is +2 and -2 respectively. Thus the oxidation number of magnesium is increases and the oxidation number of oxygen is decreased. So magnesium is oxidized and oxygen is reduced.
    Mg + O2 → 2MgO
  2. The reaction between Sodium and Chlorine: From the same way, we can explain this reaction. Here oxidation number of Sodium and Chlorine is zero before the reaction and +1 and -1 after the reaction. Thus sodium oxidized and chlorine reduced.
    Na + Cl2 → 2NaCl
Problem
    Why sulfur dioxide has properties of oxidation and reduction?
Answer
    This can be explained by the oxidation state of sulfur. The oxidation numbers of sulfur in the compounds H2S, SO2, and SO3 are -2, +4 and +6 respectively. Thus the highest oxidation state of sulfur is +6 and the lowest is -2. 
    The oxidation state of the free sulfur element is 0. In SO2, the oxidation state of sulfur is +4, this oxidation state is the middle of 0 and +6.
Thus, the oxidation state can increases from +4 to +6 and decreases from +4 to 0.

SO₂ + H₂S → 2H₂O + 3S

Definition of lewis acid and lewis base

In 1923, the American chemist Gilbert Newton Lewis proposed generalized definitions for acids that do not restrict them to compounds containing hydrogen but in terms of the structure of acids and bases.

Lewis defined an acid as any molecular or ionic species that contains an atom capable of sharing a pair of electrons furnished by a base. In the simplest terms, a Lewis acid is an electron acceptor.
Lewis base was to be considered according to Lewis as an electron pair donor and Lewis acid as an electron pair acceptor.

Difference between lewis acid and lewis base

Lewis acid, therefore, combines with a base leading to a coordinate link from the base to the acid.
An acid must have a vacant orbital into which an electron pair donated by a base can be accommodated. A base, then, is a substance which has at least one unshared (lone) pair of electrons.

H⁺ ← : NH₃

According to the lewis, the neutralization reaction is simply the formation of a coordinate bond between an acid and a base.

The neutralization product, termed as coordinate complex or adduct, may be either non-ionizable or may undergo dissociation or condensation reaction depending on its stability.
H⁺ + : NH₃ H ← : NH₃]⁺
Acid Base Adduct
F₃B + :NH₃ ⇄ F₃B ← :NH₃

What is the difference between a Bronsted and Lewis acid?

Bronsted-Lawery defined, a base is any molecules or ion, which accepts a proton. Lewis defined an acid as any molecular or ionic species that contains an atom capable of sharing a pair of electrons furnished by a base.

So that during neutralization reaction a coordinate covalent bond is formed between the donor atom of the base and the proton.

H⁺ + :NH₃ ⇄ [H ← :NH₃]⁺

Thus all Bronsted-Lawery bases are bases in Lewis sense also.
Definition of lewis acid and lewis base
Lewis concept
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Are all Arrhenius acids Lewis acids?

Arrhenius defined, an acid when dissolved in water dissociate into hydrogen ions and anions, that is it releases hydrogen ions (H⁺). Since the proton can receive electron pair from a base. Thus all the Arrhenius acids are also acids under the Lewis definition.

Solvent system and lewis concept of acid and base

The solvent system concept recognizes that ammonium chloride should behave as an acid in liquid ammonia. It may be noted that ammonium ion is but an ammoniated proton and the proton can accept a lone pair of an electron from a base. Therefore an ammonium ion in liquid ammonia is also acid in Lewis sense.

The solvent system considers amide ion is a base in liquid ammonia. The amide ion is capable of donating a lone pair to an acid, hence this is a base in Lewis sense.

The reaction between the ammonium ion and amide ion in liquid ammonia is represented in Lewis sense as follows:

NH₃ + [H⁺ + NH₂⁻] ⇄ NH₄⁺ + NH₂⁻

Question
What is the conjugate acid of bisulfate?

Answer
Bisulphate ion is HSO₃. It may lose a proton to give the conjugate base SO₃⁻², thus behaving as an acid. Again it may add on a proton to give the conjugate acid, thus showing its base character.

How to classify lewis acids and bases?


A wide variety of compounds are recognized as acids and bases according to Lewis. Any Lewis acid must contain at least one empty orbital in the valence shell of one of its atoms to accept an electron pair from a Lewis base.

An incomplete octet of the central atom

Typical examples of this class of acids are electron-deficient molecules such as alkyl and halides of Be, B, and Al. Some reactions of these types of Lewis acid with Lewis base are,

F₃B + O(C₂H₅)₂ F₃B←O(C₂H₅)₂
Acid Base Adduct
Cl₃Al + NC₅H₅ ⇄ Cl₃Al ← NC₅H₅
Me₃B + N₂H₄ ⇄ Me₃B ← N₂H₄

Question
What happens when BF₃ reacts with NH₃?

Answer
In ammonia, the central nitrogen atom has lone pair of electrons This lone pair coordinate to the empty orbital of an atom, which is termed as a base according to Lewis.

The compounds with less than an octet for the central atom are Lewis acids. In BF₃, B in BF₃ contains six electrons in the central atom thus it behaves as an acid.
F₃B + : NH₃ F₃B←NH₃
Acid Base Adduct

Elements having vacant d orbital

The central atom of the halides such as SiX₄, GeX₄, TiCl₄, SnX₄, PX₃, PF₅, SF₄, SeF₄, TeCl₄, etc. have vacant d-subshell.

These substances can, therefore, accept an electron pair from the Lewis base to accommodate in their vacant d-subshell and can thus form adducts with a number of halide ions and organic bases. These substances are thus Lewis acids.

These halides are vigorously hydrolyzed the appropriate HX. The hydrolysis reaction takes place through the formation of an unstable adduct(intermediate) with H₂O.

A cation is formed when an atom

Theoretically, all simple cations are potential Lewis acids. Reactions of some cations as Lewis acids with Lewis base are,

Lewis acid + Lewis BaseAdduct

Ag⁺ + 2(:NH₃) ⇄ [NH₃→Ag←NH₃]⁺

Cu⁺² + 2(:NH₃) ⇄ [Cu(NH3)₄]⁺²

Co⁺³ + 6(: OH₂) ⇄ [Co(OH₂)₆]⁺³

Li⁺ +: OHCH₃ ⇄ [Li←OHCH₃]
The Lewis acid strength or coordinating ability of the simple cations are increased with-
  1. An increase in the positive charge carried by the cation.
  2. An increase in the nuclear charge for atoms in any period of the periodic table.
  3. A decrease in ionic radius.
  4. A decrease in the number of the shielding shells.
Evidently, the acid strength of simple cations increases for the elements on moving from left to right in a period and from bottom to top in a group on the periodic table.

Strength of Lewis acid
Fe⁺²ㄑFe⁺³ ( Positive charge increases from +2 to +3) K⁺ㄑNa⁺ (On moving from bottom to top in a group) Li⁺ㄑBe⁺² (On moving from left to right in a period)

Molecules having multiple bonds between atoms of dissimilar electronegativity

Example of this class is CO₂, SO₂, and SO₃. In these compounds, the oxygen atom is more electronegative than sulfur or carbon atom. As a result, the electron density of π-electrons is displaced away from carbon to sulfur atoms which are less electronegative than oxygen, towards the O-atom.
Thus carbon and sulfur are electron deficient able to accept the electron pair from a Lewis base such as OH⁻ ions to form a dative bond.

Is SnCl4 a Lewis acid or base?


Bi-positive tin can function both as a Lewis acid and a Lewis base. In SnCl₂ Lewis shows a lone pair on tin through it is as yet short of an octet. Ligands, particularly donor solvents with lone pairs, may coordinate to tin giving complexes. Here SnCl₂ behaves as an acid in Lewis sense.

Again interaction of platinum group metal compounds with SnCl₂ (SnCl₃⁻) as donor leads to the formation of coordinate complexes.

[(Ph₃P)PtCl(SnCl₃)], [RuCl(SnCl₃)₂]⁻²

These are examples of the Lewis base behavior of SnCl₂.

Both silicon and tin are members of Group IV of the periodic table and their quantum shells admit of d-orbitals. As a result, they can expand their valence shell through SP3d2 hybridization and can give rise to six-coordinate complexes.

In fact complex like [SiCl₄{N(CH₃)₃}₂], [SiCl₂(Py)₄]Cl₂, [SnCl₄(Bpy)] are known. Thus SiCl₄ and SnCl₄ can function as Lewis acids.

Six electrons can exist in a p-orbital

Oxygen and sulfur contain six electrons in their valence shell and therefore regarded as Lewis acids.
The oxidation of SO₃⁻² to SO₄⁻² ion by oxygen and S₂O₃⁻² ion by sulfur is the acid-base reaction.
SO₃⁻² + [O] ⇄ [O←SO₃]⁻²
SO₃⁻² + [S] ⇄ [S←SO₃]⁻²

Question
What type of reaction are sulfur trioxide and water?

Answer
SO₃ like BF₃ has less than an octet and will be termed as an acid according to Lewis concept.
But in H₂O oxygen atom contain lone pairs to donate the Lewis acid.

How to use lewis theory to determine the chemical formula?

  1. This concept includes those reactions in which no proton is involved.
  2. Lewis's concept is more general than the Bronsted - Lowry concept or protonic concept in that acid-base behavior is not dependent on the presence of one particular element or on the presence or absence of a solvent.
  3. It explains the long-accepted basic properties of metallic oxides and the acidic properties of non-metallic oxides.
  4. This theory explains many reactions such as gas-phase high temperature and non-solvent reaction as the neutralization process.

Why PCl₃ can act both as lewis acid and lewis base?

PCl₃ has a lone pair of electrons in phosphorus. This lone pair may coordinate with a metal ion thus allowing the compound to serve as a Lewis base.

Ni + 4PCl₃ [Ni(PCl₃)₄]⁰
Acid Base Adduct

Again the quantum shell of phosphorus has provision for d-suborbital which can receive back donated electrons from electron reach low oxidation state of a metal ion. In this latter case, the tri-covalent phosphorus compound serves as a Lewis acid.

Acidic nature of oxides

Acidic oxides react with water to give oxoacids. The higher the oxidation number and the higher the electronegativity the greater the central element will force to reaction with water to give the oxoacid.

Of all these oxoacids nitrogen has the highest oxidation number and electronegativity. Thus the order of acidic nature and oxidation number are-

Acidic nature
Na₂OㄑMgOㄑAs₂O₃ㄑN₂O₅

Oxidation number of oxide-forming element
+1, +2, +3, +5.

Conjugate acid-base pair definition chemistry

A great advancement was made in the area of acids and bases when Bronsted and Lowry proposed in 1923 a new concept that is independent of solvents.

According to them, an acid is defined as any hydrogen-containing material (a molecule or a cation or an anion) that can release a proton (H⁺) to any other substance.

Whereas a base is a substance (a molecule or a cation or an anion) that can accept a proton from any other substance.
Conjugate acid-base pair definition chemistry
Bronsted Lowery concept
In short, an acid is a proton donor and a base is proton acceptor. A neutralization process has therefore involved a release of a proton by acid and the acceptance of proton by the base. Neutral compounds or even ions could be designated as acids or bases according to this concept.

What are the Bronsted acids?

HCl(molecular) → H⁺ + Cl⁻

[Al(H₂O)₆]⁺³(cationic) → H⁺ + [Al(H₂O)₅(OH)]⁺²

HCO₃⁻(anionic) → H⁺ + CO₃⁻²

What are the Bronsted bases?

Pyridine(molecular) + H⁺→ [Py(H)]⁺

[Al(H₂O)₅(OH)]⁺²(cationic) + H⁺ → [Al(H₂O)₆]⁺³

H⁺ + CO₃⁻²(anionic) → HCO₃

What are conjugate acid-base pairs give example?

When an acid release a proton (H⁺), the residue must be a base and this base can take up a proton (H⁺) to form original acid. Thus acid-base neutralization involved two acid or two bases, thus forming conjugated acid-base pairs.

Thus, a conjugate base of an acid is that part left after the proton lost. Similarly, the conjugate acid of the base is the species formed on the addition of a proton to the base.

HCl(acid) + H₂O(base) → H₃O⁺(acid) + Cl⁻(base)

In the above reaction, HCl donates a proton to H₂O and is, therefore, an acid. H₂O, on the other hand, accepts a proton from HCl and is, therefore, a base.

In the reverse reaction which at equilibrium proceeds at the same rate as the forward reaction, H₃O⁺ ion donates a proton to Cl⁻ ion, hence H₃O⁺ ion is an acid. Cl⁻ ion is a base because it accepts a proton from H₃O⁺ ion.

HCl + H₂O ⇆ H₃O⁺ + Cl⁻

The members of which can be formed from each other mutually by the gain or loss of protons are called conjugate acid-base pairs.
What are conjugate acid-base pairs give example?
Conjugate acid-base pair
A conjugate base of an acid is that part left after the proton (H⁺) is lost. Similarly, the conjugate acid of a base is the species formed on the addition of a proton to a base.

An acid exhibits its acid properties only when it is allowed to react with a base. Similarly, a base displays its basic properties only when it exposed to an acid.

Acid-base reaction equations

Acid-base reaction equations
Acid-base reaction
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How do you identify acids and bases from an equation?

Conjugate acid of a species is the one that is obtained on the addition of a proton and the conjugate base of a species is one that is obtained on the release of a proton.

H₂O(acid1) + HX⁻(base2) ⇆ OH⁻(base1) + H₂X(acid2)

In the above reaction, HX⁻ acts as a base and its conjugate acid is H₂X.
HX⁻(acid1)+ H₂O(base2) ⇆ X⁻²(base1) + H₃O⁺(acid2)

In this reaction HX⁻ acts as an acid thus its conjugated base is X⁻². In the same way, the conjugate acid of X⁻² is HX⁻ but X⁻² cannot have any conjugated base because there is no proton that can release.

Acidic character in the periodic table

These hydrides become increasingly acidic order
CH₄ㄑNH₃ㄑH₂OㄑHF.

Methane has negligible acidic properties, but ammonia donates a proton more to the strong base to form NH₂⁻, water loses a proton even more readily and hydrogen fluoride is a fairly strong acid.

The increase in the acidic properties of these hydrides is due to the fact that as we move from CH₄ to HF in the periodic table the stability of their conjugate bases increases in the order,
CH₃⁻ㄑNH₂⁻ㄑOH⁻ㄑF⁻.

Orbital electronic configuration of atoms

Orbitals of a multi-electron atom are not likely quite the same as the hydrogen atom orbitals. For practical purposes, however, the number of orbitals and there shapes in multi-electron cases may be taken the same as for the hydrogen orbitals.

In multi-electron atoms, experimental studies of spectra show that orbitals with the same value of n but different l values have different energies.

The 3S orbital is lower energy then 3P orbitals which again are of lower energy than 3d orbitals. However, orbitals belonging to a particular type ( P or d or f ) will be of equal energy (degenerate) state of an atom or an ion.

The three P orbitals or the five d orbitals originating from the same n will be degenerate. The separation of orbitals of a major energy level into sub-levels is primarily due to the interaction among the many electrons.

This interaction leads to the following relative order of the energies of each type of orbitals

1S〈2S〈2P〈3S〈3P〈4S〈3d 〈4P〈5S〈4d〈5P〈6S〈4f〈5d 〈6P〈7S〈5f〈6d

Orbital energy levels in many-electron atoms

Admittedly it is often difficult for the readers to remember the orbital energy level of an atom. A trivial but distinctly more convenient way for electronic configuration is
Orbital energy levels in many-electron atoms
Orbital energy level
The different orbitals originating from the same principal quantum number n are written in the horizontal lines.

Now inclined parallel lines are drawn through the orbitals according to the above picture. Filling up the different orbitals by electrons will follow these lines.

An element which contains 36 electrons the electronic configuration of this element according to the above diagram is,
1S² 2S² 2P⁶ 3S² 3P⁶ 4S² 3d¹⁰ 4P⁶

Explain Aufbau building up principle with an example

The question that arises now how many electrons can be accommodated per orbital. The answer to this follows from Pauli's exclusion principle.

Pauli exclusion principle electron configuration

No two electrons of the atom can have the same four quantum numbers.

This principle tells us that in each orbital maximum of two electrons can be allowed. The two electrons have the same three quantum numbers namely the same n, same l, and the same ml.

Any conflict with the Pauli Principle can now be avoided if one of the electrons has the spin quantum numbers is (+1/2) and (-1/2).

An alternative statement of Pauli's exclusion principle is,

No more than two electrons can be placed in one and the same orbital.

When an orbital contains two electrons, the electrons are paired. These two electrons per orbital are given the maximum accommodation of electrons.

How to find the number of electrons in an element?

The total number of electrons for a particular n is given by 2n².
How to find the number of electrons in an element?
Number of electrons in an element
To determine the electronic configuration of elements electrons are filling in different orbitals obeying certain rules.
Rules Aufbau or Building up the principle
  1. Electrons are fed into orbitals in order of increasing energy (increasing n ) until all the electrons have been accommodated.
  2. Electrons will tend to maintain maximum spin. So long orbitals of similar energy are available for occupation electrons will prefer to remain unpaired.
  3. In other words, electrons tend to avoid the same orbital, that is, hate to share space. This rule is known as Hund's Rule of maximum spin Multiplicity.
  4. Spin pairing can occur only when vacant orbitals of similar energy are not available for occupation, and when the next available vacant orbital is of higher energy.

Electronic configuration hydrogen to neon

Hydrogen has its only one electron in the 1S orbital and this electronic configuration represented below. the bar on the pictorial representation indicates the orbital and the arrow ↑ a single spinning electron in the orbital.

Helium, the second electron occupies the 1S orbital since the next 2S orbital is much higher energy.
Obeying Pauli principle the configuration 1S² represented below.

From the above rule, we can represent the electronic configuration of hydrogen to neon.
Hydrogen (H) 1S¹
1S
Helium (He) 1S²
↑↓
1S
Lithium (Li) 1S² 2S¹
↑↓
1S 2S
Beryllium (Be) 1S² 2S²
↑↓ ↑↓
1S 2S
Boron (B) 1S² 2S² 2P¹
↑↓ ↑↓

1S 2S 2Px 2Py 2Pz
Carbon (C) 1S² 2S² 2P²
↑↓ ↑↓
1S 2S 2Px 2Py 2Pz
Nitrogen (N) 1S² 2S² 2P3
↑↓ ↑↓
1S 2S 2Px 2Py 2Pz
Oxygen (O) 1S² 2S² 2P⁴
↑↓ ↑↓ ↑↓
1S 2S 2Px 2Py 2Pz
Fluorine(F) 1S² 2S² 2P⁵
↑↓ ↑↓ ↑↓ ↑↓
1S 2S 2Px 2Py 2Pz
Neon (Ne) 1S² 2S² 2P⁶
↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1S 2S 2Px 2Py 2Pz

Electronic configuration of sodium to argon atom

After neon, the next available orbital is 3S being followed by 3P. The orbitals are then progressively filled by electrons.

Thus the electronic configuration of sodium to argon given below.

Sodium (Na) 1S² 2S² 2P⁶ 3S¹
↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1S 2S 2Px 2Py 2Pz 3S
Magnesium (Mg) 1S² 2S² 2P⁶ 3S²
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1S 2S 2Px 2Py 2Pz 3S
Aluminum (Al) 1S² 2S² 2P⁶ 3S² 3P¹
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

1S 2S 2Px 2Py 2Pz 3S 3Px 3Py 3Pz
Silicon(Si) 1S² 2S² 2P⁶ 3S² 3P²
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1S 2S 2Px 2Py 2Pz 3S 3Px 3Py 3Pz
Phosphorus (P) 1S² 2S² 2P⁶ 3S² 3P³
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1S 2S 2Px 2Py 2Pz 3S 3Px 3Py 3Pz
Sulphur (S) 1S² 2S² 2P⁶ 3S² 3P⁴
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1S 2S 2Px 2Py 2Pz 3S 3Px 3Py 3Pz
Chlorine (Cl) 1S² 2S² 2P⁶ 3S² 3P⁵
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1S 2S 2Px 2Py 2Pz 3S 3Px 3Py 3Pz
Argon (Ar) 1S² 2S² 2P⁶ 3S² 3P⁶
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1S 2S 2Px 2Py 2Pz 3S 3Px 3Py 3Pz

Atomic structure of potassium and calcium

4S orbital, being of lower energy then the 3d, is filled. the elements involve potassium and calcium are represented as,

Potassium (K) [Ar] 4S¹

[Ar] 4S
Calcium (Ca) [Ar] 4S²

↑↓
[Ar] 4S
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Electronic configuration of d block elements

In Scandium(21) the twenty-first electron goes to the 3d orbital, the next available orbital of the higher energy. There are five 3d orbitals with the capacity of ten electrons. From scandium to Zinc these 3d orbitals are filled up.

The presence of partially filled d orbitals generates some special properties of the elements. Elements with partially filled d or f block elements are called transition elements.

Scandium (Sc) [Ar] 4S²3d¹

↑↓



[Ar] 4S 3d 3d 3d 3d 3d
Titanium (Ti) [Ar] 4S² 3d²

↑↓


[Ar] 4S 3d 3d 3d 3d 3d
Vanadium (V) [Ar] 4S² 3d³

↑↓

[Ar] 4S 3d 3d 3d 3d 3d
Chromium (Cr) [Ar] 4S² 3d⁴

↑↓
[Ar] 4S 3d 3d 3d 3d 3d

In reality, experimental studies on chromium revel their electronic configuration and better represented as

Chromium (Cr) [Ar] 4S¹ 3d⁵

[Ar] 4S 3d 3d 3d 3d 3d

This reordering of electrons is due to extra stability associated with a half-filled sub-shell.

Manganese (Mn) [Ar] 4S² 3d⁵

↑↓
[Ar] 4S 3d 3d 3d 3d 3d
Iron (Fe) [Ar] 4S² 3d⁶

↑↓ ↑↓
[Ar] 4S 3d 3d 3d 3d 3d
Cobalt (Co) [Ar] 4S² 3d⁷

↑↓ ↑↓ ↑↓
[Ar] 4S 3d 3d 3d 3d 3d
Nickel (Ni) [Ar] 4S² 3d⁸

↑↓ ↑↓ ↑↓ ↑↓
[Ar] 4S 3d 3d 3d 3d 3d
Copper (Cu) [Ar] 4S² 3d⁹

↑↓ ↑↓ ↑↓ ↑↓ ↑↓
[Ar] 4S 3d 3d 3d 3d 3d

Studies on copper also reveal that their electronic configuration and are better represented as

Copper (Cu) [Ar] 4S¹ 3d¹⁰

↑↓ ↑↓ ↑↓ ↑↓ ↑↓
[Ar] 4S 3d 3d 3d 3d 3d

This reordering of electrons is due to extra stability associated with a filled sub-shell.

Zinc (Zn) [Ar] 4S² 3d¹⁰

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
[Ar] 4S 3d 3d 3d 3d 3d

These 3d orbital are represented as dxy, dxz, dyz, dz², dx²-y².

Electronic configuration of gallium germanium atom

Gallium (Ga) [Ar] 4S² 3d¹⁰ 4P¹

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

[Ar] 4S 3d 3d 3d 3d 3d 4P 4P 4P
Germanium (Ge) [Ar] 4S² 3d¹⁰ 4P²

↑↓↑↓ ↑↓ ↑↓ ↑↓ ↑↓
[Ar] 4S 3d 3d 3d 3d 3d 4P 4P 4P
Arsenic (As) [Ar] 4S² 3d¹⁰ 4P³

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
[Ar] 4S 3d 3d 3d 3d 3d 4P 4P 4P
Selenium (Se) [Ar] 4S² 3d¹⁰ 4P⁴

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
[Ar] 4S 3d 3d 3d 3d 3d 4P 4P 4P
Bromine (Br) [Ar] 4S² 3d¹⁰ 4P⁵

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
[Ar] 4S 3d 3d 3d 3d 3d 4P 4P 4P
Krypton (Kr) [Ar] 4S² 3d¹⁰ 4P⁶

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
[Ar] 4S 3d 3d 3d 3d 3d 4P 4P 4P
In Gallium atomic number 31 and thirty-first electron goes to the 4P orbital, the next available orbital of the higher energy level. There are three 4P orbitals with the capacity of six electrons. From Gallium to Krypton these 4P orbitals are filled up.

Acids and bases questions and answers

We have already considered the thermodynamic aspects of the release of a proton by a molecule in the gas phase. The heat of the ionization of a proton is actually the proton affinity of the corresponding anion. Such proton affinity data are now available with a fair degree of accuracy to enable a good comparison of relative acid-strengths in the gas phase.

Acids and bases posses in the sense of some opposite properties, for example, acids turn blue litmus red and bases turn red litmus blue.

Which is more acidic phosphoric acid or phosphorous acid?

Oxidation state rule applied to the oxoacids of phosphorus (H₃PO₂, H₃PO₃, and H₃PO₄) it is predicted that the acidic character of these acids should be in the order

H₃PO₂ㄑH₃PO₃ㄑH₃PO₄
Experimental observation suggested the reverse order
H₃PO₂ ≥ H₃PO₃〉H₃PO₄

The experimental order can be explained when we consider the structures of this phosphorus acids
Which is more acidic phosphoric acid or phosphorous acid?
Hypophosphorous, phosphorus, and phosphoric acid
H₃PO₂ is a monobasic acid. The proton attached to oxygen has a far greater chance of dissociation than any directly bonded hydrogen.

The structure of H₃PO₂ involves one protonated oxygen and another unprotonated oxygen. H₃PO₃ is dibasic and hence has two protonated oxygens and one unprotonated oxygen. H₃PO₄ is tribasic, has three protonated oxygens and one unprotonated oxygen.

In this series, therefore, the number of unprotonated oxygen, which are the vehicles for the enhancement of acidity, is the same for all the three acids. But dissociable protons increase from one in H₃PO₂ to three in H₃PO₄.

Therefore the overall inductive effect of the unprotonated oxygen decreases from H₃PO₂ to H₃PO₄. Hence the acidity slightly falls off in the order

H₃PO₂ ≥ H₃PO₃〉H₃PO₄

What is the reaction of ammonia with an acid?

2NH₃ NH₄⁺ + NH₂⁻

KNH₂ K⁺ + NH₂⁻
Due to common ion affects the equilibrium shift to left and decrees the acidity of NH₃.

Why all alkali are bases but all bases are not alkali?

All the bases are dissolved in water to produce alkali whereas all the bases are not dissolved in water but all the alkali are dissolved in water. Thus all alkali are based but all bases are not alkali. Na₂O is a base because it dissociates in water to produce NaOH.
NaOH, KOH and Ca(OH)₂ dissolved in water to produce OH⁻ ion. Thus all these hydroxides are alkali.
NaOH ⇆ Na⁺ + OH⁻

KOH ⇆ K⁺ + OH⁻

Ca(OH)₂ ⇆ Ca⁺² + 2OH⁻
Al(OH)₃, Fe(OH)₃, Zn(OH)₂, etc does not dissolve in water but reacts with acids to produce salt and water. These are bases but not alkali.

What are acidic and basic oxides?

Acidic oxides react with water to give oxoacids. The higher the oxidation number and the higher the electronegativity the greater the central element will force to reaction with water to give the oxoacid.
Of all these oxoacids nitrogen has the highest oxidation number and electronegativity. Thus the order of acidic nature and oxidation number are
Na₂O MgO As₂O₃ N₂O₅
+1 +2 +3 +5

How do you find hydrogen ion concentration from pH?

If the pH of a solution is 4.5 and the concentration of H⁺ ion = [H⁺].
From the definition of pH
pH = - log[H⁺] = 4.5
or, log[H⁺] = - 4.5
∴ [H⁺] = 3.16 × 10⁻⁵

hydrogen and hydroxyl ion concentration and pH of a solution prepared by diluting 20 ml of 0.1M HCl.

[H⁺] = (20×0.1)/1000
= 0.002 = 2×10⁻³

[OH⁻] = (1×10⁻¹⁴)/[H⁺]
= (1×10⁻¹⁴)/(2×10⁻³) = 0.5×10⁻¹¹

pH = -log[H⁺]
= -log2×10⁻³
= 3-log2
= 2.7

What are conjugate acid base pairs give example?

Conjugate acid of a species is the one that is obtained on the addition of a proton and the conjugate base of a species is one that is obtained on the release of a proton.
H₂O + H₂X HX⁻¹ + H₃O⁺
Acid1
Base2 Base1 Acid2
In the above reaction, HX⁻ acts as a base and its conjugate acid is H₂X.
HX⁻ + H₂O X⁻² + H₃O⁺
Acid1
Base2 Base1 Acid2
In this acid-base neutralization, reaction HX⁻ is an acid thus its conjugated base is X⁻². In the same way, the conjugate acid of X⁻² is HX⁻ but X⁻² cannot have any conjugated base because there is no proton that can release.
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What is the conjugate acid of HSO₃⁻?

Bisulphate ion is HSO₃⁻. It may lose a proton to give the conjugate base SO3⁻², which is an acid. Again it may add on a proton to give the conjugate acid, thus showing its base character.

What is the nature of boron halides?

These boron halides have pi-interaction between filled p- orbitals of the halogen and empty p-orbital of boron. The effectiveness of the pi-interaction falls off with the increasing size of halogens so that it is the strongest in BI₃.

When boron halides receive an electron pair the pi-bond between boron and halogen has to be ruptured in order to make room for a coordinate bond.

H₃N→BF₃
Thus with BF₃, it will be hardest to rupture the pi-bond. Therefore the order of the Lewis acidity is
BF₃ㄑBCl₃ㄑBBr₃ㄑBI₃

What happens when BF₃ reacts with NH₃?

In NH₃, the central N atom has lone pair of electrons This lone pair coordinate to empty orbital, is termed as a base according to Lewis.

The compounds with less than an octet for the central atom are Lewis acids. B in BF₃ contains six electrons in the central atom, and it is an acid.

F₃B + :NH₃ ⇆ F₃B ← :NH₃

What type of compound are PCl3 and SnCl2?

Tri-covalent phosphorus compounds like PCl₃ have a lone pair of electrons in phosphorus. This lone pair may coordinate with a metal ion thus allowing the compound to serve as a Lewis base.

Ni + 4PCl₃ ⇆ [Ni(PCl₃)₄]⁰
Again the quantum energy level of phosphorus has provision for d-orbitals which can receive back donated electrons from electron reach low oxidation state of a metal ion. In this latter case, the tri-covalent phosphorus compound serves as a Lewis acid.

The Lewis representation of SnCl₂ shows a lone pair on tin through it is as yet short of an octet. Ligands, particularly donor solvents with lone pairs, may coordinate to tin giving complexes. Here SnCl₂ behaves as a Lewis acid.

Again interaction of platinum group metal compounds with SnCl₂ (SnCl₃⁻) as donor leads to the formation of coordinate complexes.

[(Ph₃P)₂PtCl(SnCl₃)], [RuCl₂(SnCl₃)₂]⁻²

Where SnCl₂ acts as a Lewis base.

Can SiCl₄ and SnCl₄ function as Lewis acids?

Both silicon and tin are members of Group IV of the periodic table and their quantum energy levels admit of d-orbitals. As a result, they can expand their valence shell through SP³d² hybridization and can give rise to six-coordinate complexes. In fact complex are,
[SiCl₄{N(CH₃)₃}₂], [SiCl₂(Py)₄]Cl₂, [SnCl₄(Bpy)]
Thus SiCl₄ and SnCl₄ can function as Lewis acids.

What is the difference between Lewis acid and Lewis base?

SO₃ like BF₃ has less than an octet and will be an acid according to Lewis. But in H₂O oxygen atom contain lone pairs to donate the Lewis acid. Oxygen and sulfur contain six electrons in their valence shell and therefore regarded as Lewis acids.

The oxidation of SO₃⁻² to SO₄⁻² ion by oxygen and S₂O₃⁻² ion by sulfur is the acid-base neutralization reaction.
SO₃⁻² + [O] ⇆ [O←S₂O₃]⁻²

SO₃⁻² + [S] ⇆ [S←S₂O₃]⁻²

Why hard and the soft acid-base combination is favorable?

Question
Explains why Hg(OH)2 dissolved readily in acidic aqueous solution but HgS does not?

Answer
In the case of Hg(OH)₂ and HgS, Hg is a soft acid and OH⁻ and S⁻² is hard to base and soft base respectively.

Evidently, HgS (soft acid + soft base) will be more stable than Hg(OH)₂ (soft acid + hard base).
More stability of HgS than that of Hg(OH)₂ explains why Hg(OH)₂ dissolved readily in acidic aqueous solution but HgS does not.

Question
Classify the following as soft and hard acids and bases. (i) H- (ii) Ni+4 (iii) I+ (iv) H+

Answer
  1. The hydride ion has a negative charge and is far too large in size compared to the hydrogen atom. Its electronegativity is quite low and it will be highly polarisable by virtue of its large size. Hence it is the soft base.
  2. The quadrivalent nickel has quite a high positive charge. Compared to the bivalent nickel its size will be much smaller. Its electronegativity will be very high and polarisability will below. Hence it is hard acid.
  3. Mono-positive iodine has a low positive charge and has a large size. It has a low electronegativity and high polarizability. Hence it is soft acids.
  4. H⁺ has the smallest size with a high positive charge density. It has no unshared pair of electrons in its valence shell. All these will give a high electronegativity and very low polarizability. Hence hydrogen ion is a hard acid.

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