Chut ki surgery in lab of mathematicalphysics.pptx

Chut ki surgery in lab of mathematicalphysics.pptx

1. Periodic Properties of Atoms(with reference to s and p block) Dr Nazia Tarannum

2. Slater’s Rule INTRODUCTION: In 1930, a scientist J.C. Slater proposed a set of empirical rules to understand the concept of Effective Nuclear Charge and to calculate the screening constant or shielding constant. He proposed a formula for calculation of Effective Nuclear Charge Z eff = Z – S where S is the Slater’s screening constant , Z is the Nuclear charge

3. Prior to explaining Slater’s rules, certain terms like nuclear c harge, shielding e ffect and effective n uclear c harge have to be understood. What is Nuclear Charge? It is the charge on the nucleus with which it attracts the electrons of the atom. Basically, the nuclear c harge is said to be equal to the atomic n umber i.e . the n umber of protons in an atom. It is denoted by the symbol Z. What is Shielding Effect? In case of multielectron atoms, as the orbitals are filled up, the electrons in the inner orbitals shield the electrons in the outer orbitals from the nucleus.

4. So, the electrons in the outer orbitals do not feel the full force or charge of the nucleus. Thus, the reduction of nuclear charge on the outermost electrons is called shielding e ffect or screening e ffect. Shielding effect is defined as a measure of the extent to which the intervening electrons shield the outer electrons from the nuclear charge. It is denoted by the symbol S. What is Effective Nuclear Charge ? Effective nuclear c harge is the actual charge felt by the outer electrons after taking into account shielding of the electrons. It is denoted by the symbol Z* or Z eff

5. The effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom. The term "effective" is used because the shielding effect of negatively charged electrons prevents higher orbital electrons from experiencing the full nuclear charge by the repelling effect of inner-layer electrons. The effective nuclear charge experienced by the outer shell electron is also called the core charge. It is possible to determine the strength of the nuclear charge by looking at the oxidation number of the atom .

6. In an atom with one electron, that electron experiences the full charge of the positive nucleus. In this case, the effective nuclear charge can be calculated from Coulomb's law. However, in an atom with many electrons the outer electrons are simultaneously attracted to the positive nucleus and repelled by the negatively charged electrons. The effective nuclear charge on such an electron is given by the following equation: Zeff = Z − S Z = The number of protons in the nucleus (atomic number), S = Average number of electrons between the nucleus and the electron (the number of nonvalenced electrons). Calculating Effective Nuclear Charge

7. Screening Effect or Shielding E ffect In a multielectron atom, the valence shell’s electrons are attracted to the nucleus, and these electrons are repelled by the electrons present in the inner shells. The actual force of attraction between the nucleus and the valence electrons is somewhat decreased by the repulsive forces acting in opposite directions. This decrease in the force of attraction exerted by the nucleus on the valence electrons due to the presence of electrons in the inner shells is called screening effect or shielding effect.

8. The magnitude of the screening effect depends upon the number of inner electrons higher the number of inner electrons greater shall be the value of the screening effect. The symbol σ represents the screening effect constant. Order of screening effect + S P d f S>P>d>f

9. Slater Rule 1) Write the electron configuration for the atom using the following design; (1 s )(2 s ,2 p )(3 s ,3 p ) (3 d ) (4 s ,4 p ) (4 d ) (4 f ) (5 s ,5 p ) (5d) ( 5f) (6s,6p)……. etc. 2) Any electrons to the right of the electron of interest contributes nothing towards shielding. 3) All other electrons in the same group as the electron of interest shield to an extent of 0.35 nuclear charge units irrespective of whether the electrons are in s, p, d, or f orbitals .

10. 4 ) In case of 1s electron shielding another 1s electron the screening constant value is taken to be 0.35. 5) If the electron of interest is an  s  or  p  electron: All electrons with one less value i.e. (n -1) value of the principal quantum number shield to an extent of 0.85 units of nuclear charge. All electrons with two or more less values i.e. ( n – 2, n – 3, n – 4 etc.) values of the principal quantum number shield to an extent of 1.00 units. 6) If the electron of interest is an  d  or  f  electron: All electrons to the left shield to an extent of 1.00 units of nuclear charge. 7) Sum the shielding amounts from steps 2 through 5 and subtract from the nuclear charge value to obtain the effective nuclear charge value.

11. Calculate Z *  for a valence electron in fluorine ( Z = 9 ) Electronic configuration of fluorine is 1 s 2 ,2 s 2 ,2 p 5 Grouping it according to slater’s rule : (1 s 2 )(2 s 2 ,2 p 5 ) Rule 2 does not apply; Now, one electron out of the 7 valence electrons becomes the electron of interest. The other remaining 6 valence electrons will contribute 0.35 each towards shielding. The electrons in (n – 1) orbitals i.e. 1s orbital will contribute 0.85 each towards shielding.

12. S = 0.35 x ( No. of electrons in the same shell i.e. n orbital ) + 0.85 x (No. of electrons in the (n – 1) shell ) S = 0.35 x 6 + 0.85 x 2 = 3.8 Z *  = Z – S = 9 – 3.8 = 5.2 for a valence electron.

13. APPLICATIONS OF SLATER’S RULE It provides a quantitative justification for the sequence of orbitals in the energy level diagram. It helps to explain the filling of ns - orbital (4s, 5s,6s etc. orbitals) prior to the filling of (n -1)d orbital (3d, 4d, 5d…etc.). Let us consider the case of Potassium ( Z = 19), in which the last electron is added to 4s orbital

14. The configuration of Potassium acc. to Slater is (1s 2 ) ( 2s 2 2p 6 ) ( 3s 2 3p 6 ) ( 4s 1 ) As the effective nuclear charge on electron in 4s orbital has to be calculated , the electrons in the same orbital i.e. n orbital will contribute 0.35 each, the electrons in (n – 1) orbital i.e. 3s and 3p orbitals will contribute S = 0.85 each and all the electrons in (n – 2, n – 3…. Etc. ) orbitals i.e. ( 2s, 2p, 1s) orbitals will contribute S =1.00 each. So , S = 0 x 0.35 + 8 x 0.85 + 10 x 1.00 = 16.80 Therefore, Effective Nuclear Charge Z* = Z – S = 19 – 16.80 = 2.20 Now, Let us assume that the last electron enters the 3d orbital rather than 4s orbital , Then the configuration acc. to Slater is (1s 2 ) ( 2s 2 2p 6 ) ( 3s 2 3p 6 ) ( 3d 1 )

15. Here the d electron is under Interest, so the electrons in the same orbital i.e. 3d orbital will contribute S = 0.35 each, where as the electrons in all the other orbitals i.e. (3s,3p,2s, 2p, 1s) will all contribute S = 1.00 each. So, S = 0 x 0.35 + 18 x 1.00 = 18.00 Therefore Effective Nuclear Charge Z* = Z – S = 19.00 – 18.00 = 1.00 On comparing the Effective Nuclear Charge of both 4s and 3d orbitals, we see that the 4s electron is under the influence of greater Effective Nuclear charge (Z eff = 2.20 ) as compared to 3d electron (Z eff = 1.00) in Potassium atom.

16. So, the electron in 4s orbital will be more attracted by the nucleus and will have lower energy than the 3d electron. Thus, the last electron will enter in the 4s orbital , rather than the 3d orbital in case of Potassium atom. Slater’s rule explain why 4s electrons are lost prior to 3d electrons during cation formation in case of Transition elements. It helps to explain why size of a cation is always smaller than its neutral atom. It explains why a anion is always larger than its neutral atom

17. LIMITATIONS OF SLATER’S RULE Slater grouped both s and p orbitals together for calculating effective nuclear charge, which is incorrect. This is because radial probability distribution curves show that s orbitals are more penetrating than p orbitals. So , the s orbitals should shield to a greater extent as compared to p orbital. According to Slater, all the s, p, d and f electrons present in shell or energy level lower than (n–1) shell will shield the outer n electrons with equal contribution of S=1.00 each. This is not justified as energetically different orbitals should not contribute equally. Slater’s rules are less reliable for heavier elements.

18. Atomic and Ionic Radii Down a group: Increases More electrons are being added and therefore shells are being filled. Shielding causes the valence electrons to be held less tightly. Across a period: Decreases More protons are being added and therefore the attraction between the electrons and protons is stronger.

19. Ionic Radius Radii of Cations are smaller than the parent atoms because there are more protons than electrons in the cation so the valence electrons are more strongly attracted to the nucleus. Na and Na+ Radii of Anions are larger than the parent atoms because the extra electrons result in more repulsion between the valence electrons. Cl and Cl -

20. Ionic Radius

21. Electronegativity The ability of an atom that is bonded to another atom to attract the bonding electrons towards itself.

22. The electronegativity of an element depends on a combination of two factors: 1. Atomic radius As radius of an atom increases , the bonding pair of electrons become further from the nucleus. They are therefore less attracted to the positive charge of the nucleus, resulting in a lower electronegativity. higher electronegativity lower electronegativity Electronegativity and atomic radius

23. Electronegativity, protons and shielding 2. The number of unshielded protons The greater the number of protons in a nucleus, the greater the attraction to the electrons in the covalent bond, resulting in higher electronegativity. However, full energy levels of electrons shield the electrons in the bond from the increased attraction of the greater nuclear charge, thus reducing electronegativity . greater nuclear charge increases electronegativity… …but extra shell of electrons increases shielding.

24. Electronegativity Trends: across a period Electronegativity increases across a period because: 1. The atomic radius decreases . 2. The charge on the nucleus increases without significant extra shielding. New electrons do not contribute much to shielding because they are added to the same principal energy level across the period.

25. Electronegativity trends: down a group Electronegativity decreases down a group because: 2. Although the charge on the nucleus increases, shielding also increases significantly. This is because electrons added down the group fill new principal energy levels. 1. The atomic radius increases .

26. Pauling's /Allred Rochow's Scale Allred- Rochow 's Electronegativity is a measure that determines the values of the electrostatic force exerted by the effective nuclear charge on the valence electrons. The value of the effective nuclear charges is estimated from  Slater's rules . The higher charge, the more likely it will attract electrons. Although, Slater's rule are partly empirical. So the Allred- Rochow electronegativity is no more rigid than the  Pauling Electronegativity . Electronegativity Pauling established Electronegativity as the "power" of an atom in a molecule to attract electron to itself. It is a measure of the atom's ability to attract electron to itself while the electron is still attached to another atom. The higher the values, the more likely that atom can pull electron from another atom into itself. Electronegativity correlates with bond polarity, ionization energy, electron affinity, effective nuclear charge, and atomic size.

27. Difference Between Electronegativity and Electron affinity Electronegativity refers to the ability of the atoms to attract the electrons from the other elements. Electron affinity refers to the amount of energy that is liberated whenever a molecule or a neutral atom tends to acquire an electron from the other elements. Cl  has higher electron affinity than F.  On comparing chloride ion with fluoride ion we find that electron density per unit volume in fluoride ion (F−) is more than in chloride ( Cl −) "ion. This means that coming electron in fluorine atom finds less attraction than in chlorine atom. Consequently, electron affinity of chlorine is higher than that of fluorine.

28. Ionization Enthalpy Ionization Enthalpy is the energy used for the withdrawal of an electron from its gaseous atom or ion. An atom or molecule’s first ionizing force ( Ei ) is the energy needed to remove one mole of electrons from one mole of separated gaseous atoms or ions. Ionization energy increases from left to right. Decreases from top to bottom. Ionization Enthalpy unit is KJ/mol. The first ionization energy of element A is the energy required by an atom to lose an electron and form A+ ion. A(g)→ A + (g)+ e− Similarly, the second ionization energy is; A + (g) → A ++ (g) + 2e−

29. Ionization enthalpy will always be positive, as one needs to provide a certain amount of energy to remove an electron from an atom or ion. The second frequency of ionization will always be higher than the first ionization energy. IE1<IE2<IE3

30. Electron Gain Enthalpy Electron gain enthalpy is also referred to as electron affinity although they have very slight differences. It is the measurement of the strength of the element that binds the additional electrons to the atom of that element. Electron gain enthalpy is described as the amount of energy released, when an electron is added to an isolated gaseous atom. The electron gain enthalpy is measured with the unit of KJ/ mol. In the process of electron gain enthalpy, either releasing of energy or absorption of energy takes place.

31. Electron Gain Enthalpy is represented by eg H .  It indicates the energy of the extra electron which got bound to the gaseous atom.  Electrons gain enthalpy will be more, when the amount of energy is released more in the chemical reaction. These types of reactions are both exothermic and endothermic reactions. The reaction is exothermic as it releases the energy and the reaction is endothermic as it intakes the energy, on the basis of the constituent elements. During the process, a negative gaseous ion called anion is formed. The electron gain enthalpy is measured in KJ/ mol and electron volts per atom.

32. Effective Nuclear Charge

33. Arrange these elements in increasing electron affinity: Mg, N, Na, F, Cl , O, C Why do non-metals have greater electron affinity than metal atoms? What is the difference between electronegativity and electron affinity? Arrange Be, B, N, O, F in increasing order of electron affinity.