trends in ionisation enthalpy of d-block elements part- 2

 

Ionisation enthalpy- M(g) -----> M⁺(g) + Δ_iH (ionisation enthalpy) 

·      Ionisation potential values increase in a period from left to right with increase in  nuclear charge. The increase, however, is not regular.

·      In vertical columns, i.e., groups, the ionisation potential decreases from first member to second member in most of the cases as expected,

·       however, the third member has higher value than second member. This is due to lanthanide contraction, i.e., the atomic radii of the elements of the same group of the second and third transition series are nearly the same but atomic numbers differ by 32. Thus, the outer electrons are held very firmly and the ionisation potential values are very high.   

·      On account of this, the members of 5d are inert under ordinary conditions, i.e., Pt, Au, Hg, etc., are noble metals.

·      Ionisation enthalpy depends upon

      a) Nuclear charge α I.E   α 1/size of atom     

      b) electronic configuration

  (order of ionisation enthalpy  à partially filled < half-filled < completely filled orbitals)

                                                       Fe                           Mn                     Zn

  Electronic configurationà3d⁶4s²                 3d⁵4s²           3d¹⁰4s²      

  Stability order                      less                 stable         most stable

·      I^st ionisation enthalpy of 5d series element are higher than those of 3d or 4d elements

 Ans it is due to lanthanoid contraction or the ineffective shielding of the nucleus by 4f electrons over effective nuclear charge.

·      Ist ionisation enthalpy of Mn is greater than Fe

Ans-due to presence of half-filled d-orbital in Mn(3d⁵4s²⁾ that is most stable than partially filled d-orbital present in Fe(3d⁶4s²)

·      Third ionisation energy of manganese is very high.

 Ans This is because the third electron has to be removed from a stable configuration, (Mn^{2 +} 3d⁵), i.e;, half-filled 3d-subshell.

·      The second ionisation potential values of Cu and Cr are sufficiently higher than those of neighbouring elements.

Ans This is because of the electronic configuration of Cu⁺ which is 3d¹⁰ (completely filled) and of Cr⁺ which is 3d⁵ (half-filled), i.e., for the second ionisation potential, the electron is to be removed from very stable configurations.

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Lewis symbol and rule of Lewis dot structure -why does atom combine with other atom?

 Chemical bond- 

·       A force that acts between two or more atoms to hold them together as a stable molecule. 

·       This process accompanied by decrease in energy.  

·       Decrease in potential energy (P.E.) is proportional to Strength of the bond.  

·       Therefore, molecules are more stable than atoms. 

Cause of Chemical Combination  

(A) Tendency to acquire minimum energy  

(a) When two atoms approach to each other. Nucleus of one atom attracts the electron of another atom.  

(b) If net result is attraction, the total energy of the molecule decreases and a chemical bond form. 

 (d) So, Attraction is inversely proportional 1/energy and proportional to Stability.  

(e) Bond formation is an exothermic process during formation of a molecule by combining of molecules 

(B) Tendency to acquire noble gas configuration:  

(a) Atom combines to acquire noble gas configuration.  

(b) Only outermost electrons or valence electrons i.e. ns, np and (n-1) d shells electrons participate in bond formation or chemical reaction. 

(c) Inert gas elements do not participate in bond formation, since they have stable electronic configuration as 1s² or ns²np⁶) 

Lewis Symbols  

A Lewis symbol is a convenient way to represent the valence electrons, which are shown as dots placed on the sides, top, or bottom of the symbol for the element.  

Na

Lewis Structures  

·       The Lewis structure is a representation of a molecule that shows how the valence  electrons are arranged among the atoms in the molecule. 

·       These representations are named after G. N. Lewis, who conceived the idea while lecturing to a class of general chemistry students in 1902. 

·       Lewis structures are based on achieving a noble gas electron configuration by atoms 

·       Writing Lewis structures is a trial-and-error process. 

·       The following steps are followed in constructing dot formulae for molecules and polyatomic ions:  

      (i) Write a symmetrical 'skeleton' for the molecules and poly atomic ions. 

                   a. The least electronegative element is usually taken as the central                                                                   element except Hydrogen. 

                   b. Oxygen atoms do not bond to each other except in O₂, 0 ₃, peroxide and superoxides. 

    c.       Hydrogen actually bonds to an oxygen atom and not to the central atom in ternary acids (oxyacids) ex-Nitrous acid HN02, has the skeleton H-O-N=O 

(ii) Calculate the number of electrons available in the valency shell of all the atoms (sum of the valence electrons for a molecule.) this is expressed by A  

         For negatively charged ions add to the total number of electrons equal to the charge on the anion and for positively charged ions, subtract the number of electrons equal to the charge on the cations 

Example – 

1). Sum of valence electrons for PO₄^3-  = 1 x 5 (for P atom) + 4 x 6 (for 0 atoms) + 3 (for charge)   

                      = 5 + 24 + 3  

                      = 32 electrons.

  

  2). Sum of valence electrons for NH₄⁺ ion  = 1 x 5 (for nitrogen atom) + 4 x 1 (for H                                      atoms) 1 (for positive charge)   

                                                                      = 5+4-1 

                                                                      =  8 electrons. 

  iii). Calculate the total number of electrons needed by all atoms to achieve noble gas                                       configuration.     

                                      This number is represented by N. 

                                         For example: N for  H2S04 

                                          N =2x2+8xl+8x4 =4+8+32 = 44 electrons. 

  iv).  Calculate the total number of electrons shared. This is represented by S which is equal to

                 N - A. For example: S for H2SO4 .

                                                S = N-A = 44- 32 =12 electrons. 

   v) Place the shared pair of electrons into the skeleton, using double and triple bonds only when                   necessary. 

  vi) Place the additional unshared (lone) pairs of electrons to fill the octet to every atom except                        hydrogen which can have only 2 electrons as the total comes equal to A 

 

States of matter intermolecular forces Part -1

 

States of matter

There are two types of properties shown by individual or the collection of a large number of particles

A)        1)   Individual/atomic properties- are shown by single particle or atom of a matter

Examples- atomic size, atomic/ionic radii, ionisation enthalpy, electron gain enthalpy, electronegativity, electro positivity

B)      2). Bulk properties-

                           those properties which are associated with the collection of a large number of individual particles are called bulk properties.

Example - Melting, boiling, solubility, water having wetting properties

Water exist in three states of matter as

                                                                   Solid -     Ice

                                                                  liquid – water

                                                                  Gas - water vapour

Triple point- the state at which all three phases exists simultaneously is called triple point

i.e. there are three states of matter as         a) solid   b) liquid     c) gas

on the basis of shape, volume, rigidity, fluidity, interparticle spaces, interparticle force, K.E. and Diffusion  

Mainly, the states of matter depend upon –

                                                       1). Energies of molecules

                                                       2). On the manner in which the individual molecules bind/aggregate

Note -Chemical properties of a substance do not change with the change of its physical state but the rate of chemical reaction do depend upon the physical state.

          Therefore  ,it become necessary for a chemist to know the knowledge of state of matter and the physical laws which govern the behaviour of matter in different states

Intermolecular forces-

·       The forces of attraction and repulsion between interacting particles (atoms or molecules) are called intermolecular forces.

·       These exist all matters ( solid liquid gas)  and responsible for many structural feature and physical properties of matter.

·       this term does not involved the electrostatic/ionic force and covalent bonds

·       intermolecular forces are weaker forces.

·       The attrative intermolecular forces are known as Vander Waals forces, in honour of Dutch scientists Johannes who studied the effect of these molecular forces on the real gases

Intermolecular forces  arise due to the following type of interactions:-

1). Dispersion forces or London force or instantaneous dipole – induced dipole forces

      ·            This force of attraction was firstly proposed by the German physicist Fritz London.

      ·           Dispersion forces are weaker intermolecular forces

      ·          These forces exert among the atoms(noble gases) or molecules(H₂ O₂ N₂ CO2  which are electrically symmetrical and have no dipole moment because their electronic charge cloud is symmetrically distributed.

       ·         Temporary dipoles can occur in non- polar molecules when the electrons that constantly orbit the nucleus occupy a similar location or unsymmetrical that means the charge cloud become more on one side than the other  by chance(momentarily).it is also called instantaneous dipole.

                     This instantaneous or transient dipole distorts the electron density of the other atom or molecule.

    i.e. Temporary dipoles can induce a dipole in neighbouring molecules, initiating an attraction called London force.

·       These forces are always attractive and interaction energy is proportional to 1

2. Dipole- induced dipole interaction·       This force of attraction was firstly proposed by the German physicist Fritz London.

·       Dispersion forces are weaker intermolecular forces

·       These forces exert among the atoms(noble gases) or molecules(H₂ O₂ N₂ CO2  which are electrically symmetrical and have no dipole moment because their electronic charge cloud is symmetrically distributed.

·       Temporary dipoles can occur in non- polar molecules when the electrons that constantly orbit the nucleus occupy a similar location or unsymmetrical that means the charge cloud become more on one side than the other  by chance(momentarily).it is also called instantaneous dipole.

·       This instantaneous or transient dipole distorts the electron density of the other atom or molecule.

i.e. Temporary dipoles can induce a dipole in neighbouring molecules, initiating an attraction called London force.

·       These forces are always attractive and interaction energy is proportional to 1/r⁶

2. Dipole- diploe (Keesom) force-

     ·       This interaction was first studied by Kessom in 1912.

     ·       These forces occur in molecules which have permanent electric dipole such as HCl, NH3, H2O etc.

     ·       A polar molecule has separate centre of positive and negative charges and the ends of dipoles possess “partial charge” (δ=delta=partial)

    ·       The partial charges are always less than the unit electronics charge (1.6 X 10^-19C).

    ·       These forces arise due to interaction between oppositely charged ends of the polar molecules.

    ·       The positive end of one molecule attract the negative end of the other molecule and vice versa

     ·       Greater the dipole moment of the molecules, greater is the forces of attraction.

     ·       It is stronger than London force

      ·       Interaction energy is proportional to 1/r³ (between stationary polar molecules and is proportional to 1/r⁶

   Besides dipole-dipole interaction, polar molecule can interact by London forces also.as a result the net intermolecular forces in polar molecules increase.

4. Hydrogen Bond

The d-and f-block elements - Metallic properties,M.P, Atomic size, atomisation enthalpy Part- I

 

d-block elements-

·       those elements in which last electron enters the d-orbital or (n-1) d orbital, are called d-block elements.

·       the elements which lie in- between s and p block elements in the long form of periodic table i.e elements of groups 3-12

·       d-block elements present in 4^th, 5^th 6^th 7^th period in periodic table also known as

1^st transition series-   from Sc (21) to Zn (30)

2^nd transition series – from Y (39) to Cd (48)

3^rd transition series – from La (57) Hf (72) to Hg (80).

4^th transition series – from Ac (89) Rf (104) to Uub (112).

·       General electronic configuration is (n-1) d^1-10 ns^1-2

Exception of 3d 4d 5d elements

3d-series  Sc (21)to Zn (30)]	4d-series [ Y (39) to Cd (48)]	5d-series  [from La (57) Hf (72) to Hg (80)]-
Cr - 3d54s1 Cu - 3d104s1	Nb (41) -4d45s1              Mo (42)- 4d55s1 Ru (44)- 4d75s1    Rh (45)- 4d85s1 Pd (46)- 4d105s1             Ag (47)- 4d105s1	Pt (78)- 5d96s1 Au (79)-5d106s1

·       Transition elements – have partially or incompletely filled d-orbitals in its ground state or in any one of its oxidation states.                              

          but Zn, Cd, Hg are not transition element since they have completely filled d -orbitals as

                     Zn- 3d¹⁰4s²,

                     Cd- 4d¹⁰5s² and

                     Hg- 5d¹⁰6s²

Important characteristics of transition elements

physical properties

1). Metallic properties- all transition elements show typical metallic properties as lustre, high tensile strength, malleability, ductility, good conductor of electricity  and heat

-   due to their small atomic sizes, large nuclear charges and the presence of unpaired d - electrons.

   Lustre nature or shiny surface due to presence of unpaired electron d-orbital

2). M.P of transition elements-In any row the melting points of these metals rise to a maximum at d5 except of Mn and Tc and fall regularly as the atomic number increases. It depends on

   Number of unpaired valence electrons —> increase strong bonding —> higher enthalpy of atomisation —> higher boiling point

3). Atomic size

In a given series, it decreases with increasing atomic number ,this is because the new  electron enters in the d-orbital each time ,the nuclear charge increase by unity

Note Similar behaviour has been observed in the second and third transition series.

·      In a vertical row, the atomic radii is expected to increase from top to bottom. Therefore, the atomic radii of transition metals of second series have larger values than those of the first transition series. However, the transition metals of third series except the first member, lanthanum, have nearly the same radii as metals of second transition series above them. This is due to Lanthanide contraction.

  i.e. Due to inclusion of fourteen lanthanides between lanthanum and hafnium, [there is continuous decrease in size from Ce(58) to Lu(7l)] hafnium size becomes nearly equal to the size of zirconium.

·      The atomic radii of 2^nd  and 3^rd  transition series (4d and 5d series) are greater as compared to 1^st transition series(3dse)

                          Or

          4d and 5d metals are very similar in size.

Ans -this is due to lanthanoid contraction, is a regular decrease in atomic radii by filling of 4f before 5d orbital. Since 4f orbital causes poor shielding effect(screening) than 5d.

Because 4f orbital is highly diffused orbital that’s why it shows poor shielding effect.

4.atomisation enthalpy- The required energy to convert 1 mole solid substance (metal) into atom is called atomisation enthalpy

     M(s) ----------> M(gas) +   Δ_aH =atomisation enthalpy

·      Why do the transition elements exhibit higher enthalpy of atomisation?

Ans

transition elements have large no of unpaired electrons, they have stronger interatomic interaction. Hence stronger metallic bonding between atoms resulting in higher atomisation enthalpy.

Concept-- since atomisation enthalpy depends on number of unpaired electrons which define metallic bond strength.

i.e. metallic bond strength α number of unpaired electrons

·   Zn has low atomisation enthalpy Ans – since it has no unpaired electron in its atom(3d¹⁰4s²)

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