1)Difference between Electronegativity(EN),Ionization energy(IE) and Electron affinity(EA).
| Ionization energy | Electronegativity | Electron affinit |
| IE is defined as the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to produce a cation.Eg: M(g)+IE1 M+ (g)+e – where,M(g) and M+ (g) represent gaseous atom and resultant gaseous cation respectively.2nd ionization energy: The energy required to remove a second electron from a singly charged gaseous cation. For Example: M+(g) +IE2→ M++ (g) + e– It is minimum for the alkali metals which have a single electron outside a closed shell.It always requires an absorption of energy for an electron to be removed from the outer shell of an atom. So,IE are usually endothermic. An atom has absolute value of IE.It is measured in eV/atom or Kcal/atom or KJ/molIt is property of isolated atom. | EN of an element is defined as the relative tendency of an atom in a molecule to attract a shared pair of electrons towards itself.Eg:Consider a bond between two atoms,A and . If the atoms are equally electro negative,both have the same tendency to attract the bonding pair of electrons, and so it will be found on average half way between the two atoms.To get a bond like this,A and B would usually have to be the same atom.If B is slightly more electronegative than A,B will attract the electron pair rather more than A.The ‘d + ‘ and ‘d – ‘ symbols indicate partial positive and negative charges.The most commonly used scale of EN is that developed by Linus Pauling in which the value 4.0 is assigned to fluorine,the most EN element. Mulliken Electronegativity : c = [EA+IE] /2 An atom has relative value of EN.It is unit less as it a number.It is property of bonded atom. | EA of an element is defined as the energy released when an electron is added to an isolated gaseous atom to form anion or negative ion.Eg: X(g)+e – → X–(g)+EA1 Where,X(g) and X–(g) represent gaseous atom and resultant gaseous anion respectively.For example,the first electron affinity of chlorine is -349 kJmol -1.By convention,the negative sign shows a release of energy. Second EA is positive because electron is being added to an already negatively charged ion.Thus,it requires more energy to overcome the repulsion. X–(g)+e –→X—(g) +EA2 Groups VIA and VIIA in the periodic table have the largest electron affinities values.First EA is negative because energy is released during addition of electron. An atom has absolute value of EA.It is measured in eV/atom or Kcal/atom or KJ/mole.It is property of isolated atom. |
2)Why IE of Nitrogen is greater than Oxygen?
Nitrogen has exactly half filled p-subshell i.e 1s 2 2s 2 2p 3 while oxygen has electronic configuration 1s 2 2s 2 2p 4 which is neither half filled nor completely filled. Therefore, it is difficult to remove an electron from N than from O. Thus first ionization energy of N is greater than O although nuclear charge of Oxygen is greater than that of Nitrogen.
3)EA of Be and Mg are almost zero.Why?
Be(1s 2 2s 2) and Mg(1s 2 2s 2 2p 6 3s 2) have completely filled s-orbital.They don’t have tendency to accept an extra electron. So, EA of Be and Mg are almost zero.
4)Why EA of Fluorine is less than that of chlorine?
As a result of small size of Fluorine, the inner electronic repulsion in the relatively compact 2p sub shell are comparatively large and repels the incoming electron thereby reducing the force of attraction of the nucleus towards the adding electron and hence decreasing the electron affinity.Consequently,the EA of F is less than that of chlorine.
5)Why inert gas shows highest IP ?
Inert gases have completely filled orbitals. We know that more the stable electronic configuration greater will be the ionization energy. For eg: Neon(1s 2 2s 2 2p 6),it has stable electronic configuration and as a result high ionization energy.
6)EA of halogens is high.Explain.
Halogen have valence shell electronic configuration ns 2 ns 5 .Due to this electronic configuration, these element have great tendency to accept an additional electron so they acquire noble gas(ns 2 np 6)configuration. Hence their EA is high.
7)Why EN of Gallium is higher than Aluminium?
Ga has a higher EN than Al due to the 10 extra electrons in the d orbitals, and protons in the nucleus. With a greater nuclear charge, the electrons are held more tightly to the nucleus and the size of Gallium decreases and EN increase.
8)What are the factors affecting Ionization energy?
a)Atomic size:
Ionization energy decreases with increase in atomic size. This is because larger the distance of the outer electron from the nucleus, the lesser will be the coulombic force of attraction between the nucleus and outer electron. In bigger atoms, outer electrons are held less firmly and it becomes easier to remove the electron. Thus ionization energy decrease with increase in atomic size. Therefore, IE is inversely proportional to atomic size. i.e IE α *1 /Atomic size
b)Nuclear charge:
By increasing the nuclear charge,electrons feel more nuclear attraction.Hence more ionization energy is required.Therefore,ionization energy is directly proportional to Nuclear charge. i.e IE α Nuclear charge
c)Penetration Power:
Tendency of becoming nearer to the nucleus is called penetration power.The order of penetration power of different sub-shells is s > p > d > f.Therefore,IE is directly proportional to penetration power.This is the reason that first ionization energy of B is less than Be.
d)Stability:
In stable configuration,more energy is required to release the electron as compared to non stable configuration.Therefore,Ionization energy is directly proportional to stability. Ionization energy is more of full-filled shell as compared to half-filled shell.
e)Screening & Shielding effect: Presence of other orbits between nucleus and last orbit decreases the nuclear attraction. This effect is called screening effect but electron-electron repulsion is called shielding effect which also decreases the nuclear attraction.Due to presence of these effects ionization energy decreases.
Downwards in a group ionization energy decreases due to increase in size of atoms while in a period from left to right ionization energy increases due to increase in nuclear charge.
9)What are the factors affecting Electron affinity?
a)Nuclear charge:
Effective nuclear charge is a measure of the attraction between a nucleus and its outer shell electrons.Clearly,it depends on the number of protons in the nucleus(i.e actual nuclear charge),but it also depends on the number of inner shells which screen the outer shell from the effect of the nucleus.Greater the nuclear charge,greater will be the attraction for the incoming electron and as a result larger will be the value of electron affinity.i.e EA α Nuclear charge.
b)Atomic size:
Larger the size of an atom,larger will be the distance between the nucleus and the incoming electron.Thus,smaller will be force of attraction felt by incoming electron and hence smaller will be the value of Electron affinity.In common,electron affinity reduces in going down the group and raise in going from left to right across the period.On moving down the group atomic size rises and on going from left to right in a period atomic size reduces. i.e EA α 1 /Atomic size
c)Electronic configuration:
Stable the configuration of an atom,lesser will be its tendency to accept an electron and hence lower the value of its electron affinity.Electron affinity of magnesium,beryllium, and calcium is practically zero.This is accredited due to extra stability of the fully completed s-orbitals in them.Therefore, if an atom has completely filled or half filled orbitals,its electron affinity will be less.
Variation of electron affinity:
Across the rows of the periodic table:
On moving from left to right in the periodic table,the electron affinity of the elements increases.The atomic size gets smaller as we move across the rows of the periodic table from left to right,and the nuclear charge/atomic number also increases simultaneously. This causes the electron affinity of atoms to be increased tremendously.Thus,we find that, on moving from left to right in the periodic table,the elements become more non metallic.
Down the groups of the periodic table:
The atomic size and nuclear charge of elements increases down the groups of the periodic table as the atomic number increases.But the effect of the increase in atomic size compensates the increase in nuclear charge.Thus,the attractive force of the nucleus on the outermost orbit electrons decreases and hence the electron affinity of the elements decreases.
10)What are the factor affecting the electronegativity?
A)Atomic size:
The electronegativity increases with a decrease in size of the atom.The smaller the size of an atom,greater is the tendency to attack the share paired of electron towards itself. Therefore,smaller atom have higher electronegativity values then larger atom.
B)Number of inner shell:
The atom having large number of inner shell has lower value of electronegativity.This is because the atom with more shell are bigger than the atom having less number of inner shells.
C)Types of hybridization:
The electronegativity increases as the s-character in hybrid orbitals increases.For e.g:the electronegativity in methane,ethene & ethyne is the increasing order as: Methane Ethene Ethyne sp 3 sp 2 sp 1 hybridization 25 % 33 % 50 % s-character The highest value of electronegativity of carbon atom in ethtyne account for its highly acidic hydrogen atom.Hence,carbon of ethyne has more electronegativity value than carbon of methane and ethene.
D)Ionization energy and electron affinity:
Higher Ionization energy and higher electron affinity lead to higher electronegativity. Higher value of ionization means large capacity of energy is required to remove the valence electron.This shows that the tendency of an atom to hold the shared pair of electron towards itself is high.It is due to this reason that element of group VIIA have highest electronegativity.
E)Charge on the ion:
A cation is the smaller in size of corresponding atom.As a result,a cation attract electron more readily than its parents atom.Thus,a cation has higher electronegativity than its parents atom.Since,the electron accepting tendency increases with an increasing in the charge on the cation,hence the cation having higher positive charge is more electronegative.Size of anion is larger as compared to that atom is less than that of its parents atom.So electron attracting tendency of an anion is less than that of its parent atom.The electronegativity of F(EN=0.8) is less than that of F atom (EN = 4).
F)Number and nature of atoms bonded to it:
The electronegativity of an atom depends upon the number and nature of the atom bonded to it.For e.g.the electronegative of phosphorous atom in PCl3 is higher than in PF5.Since F is more electronegative than Cl.Also, In PCl3 and PCl5,p of PCl5 has electronegativity value more as oxidation state of P is higher than P of PCl5.
Note:
A) Effective nuclear charge increases across each period.
B) IE,EA,EN generally increase across the periods.
C) IE,EA,EN generally decrease across the periods.
D) Removal of electrons from half filled and full filled shells requires higher energy.
E) Metals have low IE and non metals have high IE.
F) The elements with a higher IE have higher EA also.
G) Metals have low EN whereas and nonmetals have high EN.
H) Fluorine is most electronegative element.
I) First EA is exoergic while second and subsequent EA are endoergic innature.
11)S block element:
In the s-block elements the last electron enters the outermost s-orbital and as the s-orbital can accommodate only two electrons,that is why only two groups i.e group 1 & 2 belong to the s-block of the periodic table.Thus the outermost orbital of s block elements consists of one or two electrons and the orbital next to the outermost shell i.e the penultimate shell has either 2 or 8 electrons.This is the reason why the s-block elements show a fixed valency which depends on the number of electrons present in the outermost shell.
a)Electronic Configuration:
The general electronic configuration of s-block elements is ns 1 for Alkali metals and ns 2 for alkaline earth metals where n = 2 to 7.All the alkali metals have one valence electron and these loosely held s-electron in the outermost valence shell of these elements makes them the most electropositive metals.The alkaline earth metals have two electrons in the s-orbital of the valence shell.Like alkali metals,these elements are also very electropositive.
b) Metallic character:
All the Alkali metals are silvery white, soft and light metals.The Alkaline Earth metals,in general are silvery white,lustrous and relatively soft but harder than the Alkali metals. Beryllium and magnesium appear to be somewhat greyish.The Metallic character increases as we go down both groups.Both the alkali metals and the alkaline earth metals are highly malleable and ductile and have a very high tendency to lose electrons to form positive ions and hence they are highly electropositive.
c)Atomic Density:
The Alkali metals and the Alkaline Earth metals both have low density.This is because they have large ionic size due to which their atomic nuclei are widely separated in their crystal lattices.The density increases down both the groups and periods.
d)Melting and Boiling Points:
The melting and boiling points of the alkali metals are low indicating weak metallic bonding due to the presence of only a single valence electron in them.The melting and boiling points of alkaline Earth metals are higher than the corresponding alkali metals due to their smaller sizes.The trend is however not systematic but it slightly decreases down the group.
e)Oxidation state:
The Alkali metals show only +1 oxidation state,while alkaline Earth metals show +2 oxidation state only.Because of their low ionization energies, they easily lose the outermost s electron to form the uni positive ions.Once they lose the first electron,in case of alkali metals,they achieve the noble gas configuration.The same is true with respect to alkaline earth metals,when they lose the 2 electrons present in the valence shell.
f)Atomic and Ionic radii:
Both the Alkali metals and the alkaline earth metals have large atomic and ionicradii. The Atomic and the Ionic radii increases as we move down both groups.But as we go from group I to group II in the same period the atomic and the ionic radii decreases.
g)Electrode potential:
The alkali metals are strong reducing agents.The standard electrode potentials of all alkali metals lie between -2.7V and -3.0V,indicating a strong tendency to form cations in solution.The alkaline earth metals also have negative values of their standard electrode potentials.
h)Ionization Energies:
The ionization energies of the alkali metals are considerably low and decreases down the group from Lithium (Li) to Cesium(Cs).This is due to the increasing size,increasing nuclear charge, and the outermost electron is very well screened from the nuclear charge. The alkaline earth metals have low ionization energies due to their large size of the atoms.Since the atomic size increases down the group,their ionization energies decreases down the group.But as we go from group I to group II in the same period ionization energies increases.The first ionization energies of the alkaline earth metals are higher than those of the corresponding Group I metals.This is due to their small size as compared to the corresponding alkali metals.But the second ionization energies of the alkaline earth metals are smaller than those of the corresponding alkali metals.
Na → Na + + e –
Ca→ Ca ++ +2e –
i)Magnetic Properties:
Alkali metals are attracted by the applied magnetic field and hence are paramagnetic in nature whereas the alkaline earth metals are repelled by the magnetic field and hence are diamagnetic in nature.
J)Complex Formation:
Both the alkali metals and the alkaline earth metals show weak tendency to form complexes because they have no low energy vacant orbital available for bonding with lone pair of ligands.This is due to large size,low nuclear charge and poor ability to attract electrons.
K)Flame colouration:
The alkali metals and their salts,when introduced into the flame,give characteristic color to the flame.
| Li | Na | K | Rb | Cs |
| Crimson red | Golden yellow | Pale violet(Lilac) | Red-Violet | Blue |
This property of the alkali metals offers a very sensitive and reliable test for alkali metals. This property is due to the ease of excitation of the valence electrons.When elements or their compounds are introduced to flame,the electrons absorbs energy from the flame and gets excited to higher energy levels.When these electrons return to their ground state,they emit absorbed energy in form of visible light having characteristic wavelengths. Depending upon the wavelength of light emitted, different colors are imparted to the flame.Salts (generally chlorides)impart characteristic colors to the Bunsen flame.Be and Mg does not impart characteristic color to the flame.
12)P block element:
Elements belonging to the group 13 ( i.e group IIIA) to group 17 (i.e group VIIA) of the periodic table along with the group 18 i.e the zero group elements together form the p-block of the periodic table.In the p-block elements the last electron enters the outermost p orbital.They have 3 to 8 electrons in the outermost shell.As we know that the number of p orbitals is three and, therefore, the maximum number of electrons that can be accommodated in a set of p orbitals is six.Hence there are six groups of p-block elements in the periodic table numbering from 13 to 18.The First group of the p-block i.e the group IIIA is commonly called as Boron group,the second group i.e the group IVA is called Carbon group,the third group i.e the group VA is called Nitrogen group,the fourth group i.e the group VIA is called Chalcogens,the fifth group i.e the group VIIA is called Halogens and the sixth group i.e the zero group or group 18 is called Inert or Noble gases group.In the p-block all the three types of elements are present i.e.The metals,nonmetals and metalloids .
Characteristic Properties of elements in p-block of Modern Periodic Table :
a)Electronic Configuration:
The general valence shell electronic configuration of p-block elements is ns 2 np 1-6 (except for He).The inner core of the electronic configuration may,however,differ.The General electronic configuration shown by elements from group13 to 18 of p-block is as given below :
Group 13 (Boron family) – ns 2 np 1
Group 14 (Carbon family) – ns 2 np 2
Group 15 (Nitrogen family) – ns 2 np 3
Group 16 (Oxygen family) – ns 2 np 4
Group 17 (Halogen family) – ns 2 np 5
Group 18 (Noble gases) -ns 2 np 6 (except Helium)
The general electronic configuration of Helium is 1s 2 .Due to their distinct electronic configuration p-block elements show a lot of variation in properties.
b)Metallic Character:
As stated earlier p-block contains all types of elements i.e metals,non-metals and metalloids.The p-block is the only region of the periodic table to contain metalloids.The non metallic character decreases down the group whereas there is a gradual increase in non-metallic character from left to right in the p-block.The metallic character tends to increase down each group whereas it decreases as we go from left to right across a period. In fact,the heaviest element in each p-block group is the most metallic in nature.Metallic character increases as we move down the group,from B to Tl.Boron is a non metal,while all the others are metallic and good conductors of electricity.
c)Atomic Density:
The Atomic Density of elements in p-block increases down the group,this is due to increase in the size of the atom down the group.Whereas it decreases as we move from left to right across the period,this is due to the decrease in atomic size of all elements in the p- block across the period.Of all the elements,aluminium is of very low density and is widely used as a structural material. d)Melting and Boiling points:
Melting points decreases from Barium to gallium and then increases up to Tl.The low melting point of Ga is explained on the basis that Ga is diatomic in nature.Heat of sublimation and boiling points show a steady decrease.The melting and boiling points gradually increases down the group because the molecular mass increases down the group and hence the intermolecular forces increases. e)Oxidation state:
The p-block elements show a variable oxidation state.The oxidation states increases as we move from left to right in the periodic table.The maximum oxidation state shown by a p-block element is equal to the total number of valence electrons.According to this,the oxidation states shown by different groups is as follows :
Boron family (Group 13) : +3 , Carbon family (Group 14) :+4
Nitrogen family (Group 15) :+5 , Oxygen family (Group 16) : +6
Halogen family (Group 17) :+7 , Noble gases ( Group 18) : +8
Since Boron is very small, it has high ionization energy and does not lose all three valence shell electrons.So, it cannot form B3+ ion.Aluminium forms +3 ion,while Ga,In and Tl show +1 and +3 oxidation states.
f)Atomic and Ionic radii:
As we move down the group in the p-block one extra shell than the preceding element gets added into the next element.This ultimately increases the atomic and the ionic radius of every next element down the group ,which finally shows that the atomic and the ionic radii increases down the group.The trend is not same across the period.As we move from left to right in a period the atomic radii and the Ionic radii of p-block elements decreases . The atomic radius increases greatly from Boron to Aluminium.This increase is due to greater screening effect caused by the eight electrons present in the penultimate shell.
g)Electrode Potential:
The p-block elements generally have a positive electrode potential.It generally decreases down the groups.
h)Ionization Energies:
The p-block elements have high ionization potentials.The ionization energies of p-block elements increases from left to right in a period due to increasing effective nuclear charge.According to the general trends the ionization energy values decreases down the group but do not decrease smoothly as expected.Non-metals have high ionization energies than metals.It is maximum for a noble gas because noble gases have completely filled configuration.Some elements at the bottom of a group like Lead,Tin,Thallium, Bismuth etc behaves almost as a metal with very low ionization energies.
i) Magnetic Properties:
The elements Radon,Astatine,Iodine and Polonium of the p-block are non-magnetic in nature.The element Tin is paramagnetic and the rest all elements of the p-block are diamagnetic in nature.
j) Complex Formation:
The smaller size and the greater charge of the elements of different groups of p-block enable them to have a greater tendency to form complexes than the s-block elements.The complex formation tendency decreases down the group as the size of the atoms increases down the group.
k)Chemical Reactivity:
The chemical reactivity of elements in the p-block increases as we move from left to right in a period. But as we move down in a group the chemical reactivity of elements decreases down the group.
l)Reactivity of Noble gases:
All the orbitals of the noble gases are completely filled by electrons and it is very difficult to break their stability by the addition or removal of electrons.Thus the noble gases exhibit very low chemical reactivity.Because of their low reactivity noble gases,are often used when an nonreactive atmosphere is needed,such as in welding.
m)Conductivity: The conductivity of elements in p-block increases down the group.Generally the metals in the p-block are good conductors of heat and electricity whereas the non-metals are poor conductors of heat and electricity.The conductivity of metalloids lies in between the metals and non-metals.
n)Electron Negativity:
EN increases from left to right in a period and decreases down the group.On account of high values of EN,they usually form covalent compounds.
13)General Characteristics of Transition Elements:
The elements that lie in between S-block and P-block are the d-block elements.These elements are called transition elements as they show transitional properties between ‘s’ and p-block elements.These elements contain partially filled d-orbitals and hence they are called as d-block elements.The general electronic configuration of d-block elements is (n-1)d 1-10 ns 1-2 .
A)Electronic Configuration:
The general electronic configuration of d-block elements is (n-1)d 1-10 ns 1-2 .All the d-block elements except zinc,cadmium and mercury have partially filled d-orbitals.But,zinc, cadmium and mercury have completely filled d-orbitals and they exhibit common oxidation state.So,they do not come under transition elements but are studied along with d-block elements.In all the other transition elements the last electron enters the (n-1)d orbital which is called the penultimate shell.
B)Variable oxidation states:
By the study of electronic configuration of transition metals it is understood that variable oxidation state can be formed as there are both ns and (n-1)d electrons in bonding.The participation of ns electrons in bonding leads to +2 oxidation state which is a lower oxidation state.The participation of (n-1)d electrons in bonding leads to higher oxidation states like +3,+4,+5,+6 etc.These oxidation states depend upon the nature of combination of transition metals with other elements.The oxidation state increases with atomic number.This increase is related to groups.The most common oxidation state of the elements of first transition series is +2.Ionic bonds are formed in lower oxidation state transition elements whereas covalent bonds are formed in higher oxidation states. C)Magnetic properties:
By the study if electronic configuration of transition metals it is understood that they generally contain one or more unpaired electrons in the (n-1)d orbital.Due to these unpaired electrons they behave as paramagnetic substances.These substances are attracted by the magnetic field.The transition elements that contain paired electrons behave as diamagnetic substances.These substances are repelled by the magnetic field.The para- magnetic character increases as the number of unpaired electrons increases.Formation of colored compounds.Most of the transition elements form colored compounds both in solid state as well as in aqueous solution.It is already studied that the transition metals have incomplete d-orbital.The electrons are to be promoted from a lower energy level to a higher energy level.Some amount of energy is required for this process and the radiations of light are observed in the visible region.The compounds absorb a particular color from the radiation and the remaining ones are emitted. For e.g.Cu ++ are bluish green in color due to absorption of red light wavelength.As Zn has completely filled d-orbitals it cannot absorb radiation and hence Zn ++ salts are white.
D)Formation of complexes:
Transition metals form many complex ions.They are the electrically charged complexes with a metal ion in the centre which is surrounded and linked by a number of neutral molecules or negative ions.These neutral molecules or negative ions are called as ligands. As the transitions metals are small in size they form large number of complexes. For Eg: Fe(CN)6] 4-, [Cu(NH3)4] 2+ , [Ni(CN)4] —. General trends in the chemistry of first row transition series:
A)Metallic character:
Most of the transition elements of the first row form metallic bonds due to the presence of incomplete outermost energy level.So,all the transition elements exhibit metallic characters.The strength of the metallic bond depends upon the number of unpaired d-electrons.As the number increases the strength also increases.Due to the absence of unpaired electrons ‘Zn’ is not a hard metal.
B)Ionization energy:
The ionization energies of first row elements gradually increases with increase in atomic number.The ionization energy of Zn is very high than all the other metals which is dueto its fully filled d-orbital.The third ionization energy of Mn is very high than the others.
C)Oxidation State:
The first row transition elements show variable oxidation states.Zn is an exception among them.As it has fully filled d-orbital,it exhibits only +2 oxidation state.
D)Ionic radii:
In the first row transition elements the ionic radii decreases with increase in atomic number.The value of ionic radii also depends on the oxidation state of metals.As the oxidation state increases the ionic radii decreases and as the oxidation state decreases the ionic radii increases.
E)Catalytic property:
The first row transition elements exhibit catalytic properties due to the presence of unpaired electrons which can form complexes.Iron and vanadium are the most important catalysts.Iron is used as catalyst in the manufacture of ammonia.Vanadium is used in the form of vanadium pentoxide in the manufacture of sulphuric acid.
F)Coloured ions:
In the first row transition elements all the elements except Zn form colored ions.As these elements have incomplete d-orbital,some amount of energy is required to promote the electrons from lower energy level to higher energy level.This process exhibits radiations from which the compounds absorb a particular color.But some elements other than Zn also appear colorless depending on their oxidation state.
G)Alloy formation:
When one metal mixes up with another metal alloys are formed.As the d-block elements have same atomic sizes they can easily take up positions of one another.This causes alloy formation.For example:Cr,V, Mn are used in formation of alloy steels.
14)Why Cu, Zn, Hg are not considered as transition metal?
Cu=atomic number =29=3d 10 4s 1
Zn=atomic number =30=3d 104s 2
Hg=atomic number =80=5d 10 6s 2
As to be a transitional element,atom should have incomplete d-orbitals,these three elements have completely filled d-orbitals i.e having (n-1)d 10 ns 2 configuration.so they cannot be regarded as true transition element.
15)Transitional element exhibit variable oxidation states.Explain.
=The variable oxidation state of transitional element is due to availability of both (n-1)d and ns electron for bond formation,as the energies of d-orbital and s orbital are nearly equal.Most of the transitional element has two ns electron,so generally exhibit +2 oxidation state.In addition,they may also utilize one more (n-1)d electron for bond formation.So oxidation state of +3 and higher will be occur.The sum of ns and unpaired (n-1)d electron determines the highest oxidation state shown by Transitional element.
16)Why Zn ++ salts are colourless?
=The colour of transition metal ions is associated with incompletely filled (n-1)d orbitals. The transitional metals ions containing unpaired d-electron undergo electronic transitions from one d orbital to another d-orbital.During d-d transition process,they absorb certain radiation from visible light and emit the remainder as coloured light.But,in Zn ++ (3d 10 4s 0) has no unpaired d electron,so it’s salts are white.
17)Why transitional element form significant number of complexes? =Transitional element forms significant number of complexes because transitional metals yields small,highly charged ions,which have vacant (n-1)d orbitals of approximately the appropriate energy to accept the pairs of electrons donated by other group or molecules. The molecules or ions which add themselves to the cation of transitional metals are called ligands. For eg:[Fe(CN)6] 4- and [Fe(CN)6] 3-
18)Why Transitional elements are mostly para magnetic?
=Transitional elements are mostly paramagnetic because there is presence of unpaired electron in the (n-1)d orbital or (n-2)f orbital.Thus only those atoms or ions having unpaired electron shows para magnetism.The greater the number of unpaired electron in a atom or ions,the more strongly paramagnetic it is. For Eg:Fe +++ is more strongly paramagnetic than Fe ++ because Fe +++ possesses greater number of unpaired electrons.
19)Why is Cu(I) dimagnetic while Cu(II) is paramagnetic?
=Cu + (3d 10 )is dimagnetic due to absence of unpaired d-electron whereas Cu ++ (3d 9) is paramagnetic due to presence of one unpaired d-electron.
20)Why are transitional elements metals?
=Transitional elements are called metals because they have low IE,they lose electrons to form cations and exhibit metallic character.
21)Define Transitional element.
=Transitional elements are those elements whose atom or atleast one of its ions has incompletely filled d or f orbitals. Eg:Sc,Fe,Cr,Pd.
22)Define and explain inert pair effect.Clarify the reason of inert pair effectwith suitable examples. =The inert pair effect is the tendency of the electrons in the outermost atomic s orbital to remain unionized or unshared in compounds of post-transition metals.The term inert pair effect is often used in relation to the increasing stability of oxidation states that are 2 less than the group valency for the heavier elements of groups 13,14,15 and 16. In the elements of 4th,5th and 6th period of p-block elements which come after d-block elements the electrons present in intervening d and f-orbitals do not shield the s-electrons of the valence shell effectively as a result ns 2 electrons remain more tightly held by nucleus and hence do not participate in bonding.This is called inert pair effect. In other words tendency of s-electrons of valence shell to participate in bond-formation decreases.Another reason for inert pair effect is that as the size of atoms increases down the group in p-block elements the energy required to unpair the ns 2 electrons is not compensated by the energy released in forming two additional bonds.So the bond formation by valence ‘s’ electrons is not energetically favorable.The inert pair effect becomes more predominant down the group in p-block elements because of increased nuclear charge which outweighs the effect of corresponding increase in atomic size.The s-electrons thus become more tightly held and therefore become more reluctant to participate in bond formation.
For example:
1)In 3rd A group,thallium(1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 6 4f 14 5d 10 6s 2 6p 1)can exhibit +1 and+3 oxidation states but it is stable in +1 oxidation state only due to inert pair effect.
2) In 4th A group, lead([Xe] 4f 14 5d 10 6s 2 6p 2)shows both +2 and +4 oxidation states but it is stable in +2 oxidation state due to inert pair effect.
Reference 1 :
S.Chand PublishingA Textbook of Engineering Chemistry, S.Chand And Company Limited