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

Lewis concept of acids and bases
Lewis concept
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.
    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

Lewis vs Bronsted-Lawry acids and bases

    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.

Arrhenius vs Lewis concept of acids and bases

    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 vs Lewis concept

    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⁻

Classification of 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.

A central atom with an incomplete octet

    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

A central atom with vacant d-orbitals

    The central atom of the halides such as SiX₄, GeX₄, TiCl₄, SnX₄, PX₃, PF₅, SF₄, SeF₄, TeCl₄, etc. have vacant d-orbitals.
    These substances can, therefore, accept an electron pair from the Lewis base to accommodate in their vacant d-orbital 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.

The simple cation of 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 increasing
    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.

Element with an electron sextet

    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

The utility of Lewis Concept

  1. This concept includes those reactions in which no proton is involved.
  2. Lewis's concept is more general than the Bronsted - Lowry concept (that is 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.

Questions answers

Question
    Explain - ammonia behaves as a base but boron trifluride as an Acid.
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
Question
    Explain why tri-covalent phosphorus compounds can serve both as Lewis acids and also as bases?
Answer
    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
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.
Question
    Bi-positive tin can function both as a Lewis acid and a Lewis base. explain?
Answer
    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. As for examples,
    [(Ph₃P)PtCl(SnCl₃)], [RuCl(SnCl₃)₂]⁻²
    These are examples of the Lewis base behavior of SnCl₂.
Question
    Can SiCl₄ and SnCl₄ function as Lewis acids?
Answer
    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.
Question
    Arrange these oxides in order of their acidic nature: N₂O₅, As₂O₃, Na₂O, MgO.
Answer
    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.
Question
    Bisulphate ion can be viewed both as an acid and a base. explain.
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.
Question
    SO₃ behaves as acid and H₂O as a base. Explain.
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.

    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.
Bronsted Lowery concept of conjugate acid base
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.

Bronsted acids

HCl(molecular) → H⁺ + Cl⁻

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

HCO₃⁻(anionic) → H⁺ + CO₃⁻² 

Bronsted bases

Pyridine(molecular) + H⁺→ [Py(H)]⁺

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

H⁺ + CO₃⁻²(anionic) → HCO₃

Conjugate acid-base pairs

    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.
Conjugate acid-base pair
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 reactions

Acid-base reaction and conjugate acid base pair
Acid-base reaction
Question
    Name the conjugate acids and the conjugate bases of HX⁻ and X.
Answer
    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.
Question
    Arrange the following compounds in order of increasing acid strength: (i) NH₃ (ii) CH₄ (iii) HF (iv) H₂O
Answer
    These hydrides become increasingly acidic order as,
CH₄ㄑNH₃ㄑH₂OㄑHF
    Thus CH₄ has negligible acidic properties, but NH₃ donates a proton more to the strong base to form NH₂⁻, H₂O loses a proton even more readily and HF 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, the stability of their conjugate bases increases in the order,
CH₃⁻ㄑNH₂⁻ㄑOH⁻ㄑF⁻

Orbitals of a multi-electron atom

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.

For example, 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 Occupancy Order

Admittedly it is often difficult for the readers to remember the orbital energy diagram. A trivial but distinctly more convenient way for electronic configuration is
Electronic Configuration of first thirty six Elements with Pictorial Representation.
Orbital Occupancy Order
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

Aufbau or Building Up Principle

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

Pauli's exclusion principle

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).

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 contain two electrons, the electrons are paired. These two electrons per orbital are given the maximum accommodation of electrons.

Capacities of electronic levels

The total number of electrons for a particular n is given by 2n².

Capacity of electronic level
Capacity of electronic level
To determine the electronic configuration of elements electrons are filling in different orbitals obeying certain rules.
Rules Aufbau or Building up 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 of elements

Hydrogen to helium and lithium 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

Electron distribution from sodium to argon

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

Elements potassium and Calcium

Then 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

3d-block elements scandium to zinc

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

Experimental 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.

Distribution from gallium to krypton

In Gallium(31) the thirty-first electron goes to the 4P orbital, the next available orbital of the higher energy. There are three 4P orbitals with the capacity of six electrons. From Gallium to Krypton these 4P orbitals are filled up.

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

Properties of acids and bases

Some chemical substances as acids and certain others as bases rest on some experimental facts about these substances. Acids and bases posses in the sense of some opposite properties.

For example, acids turn blue litmus red and bases turn red litmus blue. On this page, the different concepts relate to questions answers we can discuss.

Neutralization of acids and bases

Question
What is the acidic order of oxoacids of phosphorus?

Answer
When the 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
But the experimental observation suggested the reverse order
H₃PO₂ ≥ H₃PO₃〉H₃PO₄

The experimental order can be explained when we consider the structures of these acids given as,

Acids and bases questions and answers
Oxoacids of phosphorus
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

Acid-base neutralization

Question
What will be the effect of adding KNH₂ to liquid ammonia with respect to acidity?

Answer
2NH₃ NH₄⁺ + NH₂⁻

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

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

Answer
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. Thus These are bases but not alkali.

Question
Arrange these oxides in order of their acidic nature: N₂O₅, As₂O₃, Na₂O, MgO.

Answer
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

What is ph in chemistry?

Question
The pH of a solution is 4.5. Calculate the concentration of H⁺ ion.
Answer
We have from the definition,
pH = - log[H⁺] = 4.5
or, log[H⁺] = - 4.5
∴ [H⁺] = 3.16 × 10⁻⁵

Question
Calculate the [H⁺], [OH⁻] and pH of a solution prepared by diluting 20 ml of 0.1M HCl to one lit.
Answer
[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
Question
Calculate the [H⁺], [OH⁻] and pH of a solution obtained by dissolving 28 gm KOH to make 200 ml of a solution.

Answer
[OH⁻] = (28×1000)/(56×200) = 2.5 M
(Molecular Weight of KOH = 56gm)

[H⁺] = (1×10⁻¹⁴)/[OH⁻] = (1×10⁻¹⁴)/(2.5)
= 4×10⁻¹⁵pH = -log[H⁺]
= - log4 × 10⁻¹⁵
= (15 - log4)

Conjugate acid base pairs

Question
Name the conjugate acid-base pairs of HX⁻¹ and X⁻².

Answer
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.

Question
Bisulphate ion can be viewed both acid-base, explain.

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

Lewis concept of acids and bases

Question
Arrange in order of Lewis acidity : (i) BF₃ (ii) BCl₃ (iii) BI₃ (iv) BBr₃

Answer
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

Question
Explain - NH₃ behaves as a base but BF₃ is acid.


Answer
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 concept.

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
Question
Explain why tri-covalent phosphorus compounds can serve both as Lewis acids and also as bases?

Answer
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 shell 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.

Question
Bi-positive tin can function both as a Lewis acid and a Lewis base. Explain?

Answer
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. As for examples,

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

These are examples, where SnCl2 acts as a Lewis base.

Question
Can SiCl₄ and SnCl₄ function as Lewis acids?

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

Question
SO3 Behaves as acid and H2O as a base. Explain.

Answer
SO₃ like BF₃ has less than an octet and will be an acid according to Lewis concept. 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

Soft and hard acids and bases

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 H⁺ is a hard acid.

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