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Atomic Structure and Periodicity

Atomic Structure and Periodicity. Schrodinger Wave Equation. Equation for probability of a single electron being found along a single axis (x-axis). Erwin Schrodinger. Pauli Exclusion Principle. No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli.

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Atomic Structure and Periodicity

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  1. Atomic Structure and Periodicity

  2. Schrodinger Wave Equation Equation for probability of a single electron being found along a single axis (x-axis) Erwin Schrodinger

  3. Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli

  4. Heisenberg Uncertainty Principle “One cannot simultaneously determine both the position and momentum of an electron.” You can find out where the electron is, but not where it is going. OR… You can find out where the electron is going, but not where it is! Werner Heisenberg

  5. Quantum Numbers Each electron in an atom has a unique set of 4 quantum numbers which describe it. • Principal quantum number • Angular momentum quantum number • Magnetic quantum number • Spin quantum number

  6. Principal Quantum Number Generally symbolized by n, it denotes the shell (energy level) in which the electron is located. Tell the energy Number of electrons that can fit in a shell: 2n2 N= 1, 2, 3, ….

  7. Angular Momentum Quantum Number The angular momentum quantum number, generally symbolized by l, denotes the orbital (subshell) in which the electron is located. n2= the number oforbitals on an energy level All orbitals with the same value for N are said to be degenerate.

  8. Magnetic Quantum Number The magnetic quantum number, generally symbolized by m, denotes the orientation of the electron’s orbital with respect to the three axes in space.

  9. Assigning the Numbers • The three quantum numbers (n, l, and m) are integers. • The principal quantum number (n) cannot be zero. • n must be 1, 2, 3, etc. • The angular momentum quantum number (l ) can be any integer between 0 and n - 1. • For n = 3, lcan be either 0, 1, or 2. • The magnetic quantum number (ml) can be any integer between -l and +l. • For l = 2, m can be either -2, -1, 0, +1, +2.

  10. Principle, angular momentum, and magnetic quantum numbers: n, l, and ml

  11. Spin Quantum Number Spin quantum number denotes the behavior (direction of spin) of an electron within a magnetic field. Possibilities for electron spin:

  12. Magnetic Properties • Although an electron behaves like a tiny magnet, two electrons that are opposite in spin cancel each other. Only atoms with unpaired electrons exhibit magnetic susceptibility • A paramagnetic substance is one that is weakly attracted by a magnetic field, usually the result of unpaired electrons. • A diamagnetic substance is not attracted by a magnetic field generally because it has only paired electrons.

  13. Determination of Atomic Radius Half of the distance between nucli in covalently bonded diatomic molecule "covalent atomic radii" Periodic Trends in Atomic Radius • Radius decreases across a period Increased effective nuclear charge due to decreased shielding • Radius increases down a group Addition of principal quantum levels

  14. Bond Radius

  15. Table of Atomic Radii

  16. I.E. ↑ n constant r ↓ Zeff↑ Zeff constant n ↑ r ↑ I.E. ↓ II/

  17. Ionization Energy: the energy required to remove an electron from an atom • Tend to be positive values • endothermic process • A + energy = A+ + e-

  18. HIGHER IONIZATION ENERGIES

  19. big jump in I.E. when core electrons start to be removed: • electrons from a lower main shell start to get removed. II/

  20. Ionization Energy: the energy required to remove an electron from an atom • Increases for successive electrons taken from • the same atom • Tends to increase across a period Electrons in the same quantum level do not shield as effectively as electrons in inner levels Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove • Tends to decrease down a group Outer electrons are farther from the nucleus

  21. Trends in Electron Affinity • The first occurs between Groups IA and IIA. • Added electron must go in p-orbital, not s-orbital. • Electron is farther from nucleus and feels repulsion from s-electrons.

  22. Trends in Electron Affinity A + e- + energy → A- A + e- → A- + energy • The second occurs between Groups IVA and VA. • Group VA has no empty orbitals. • Extra electron must go into occupied orbital, creating repulsion.

  23. Trends in Electron Affinity .

  24. Electron Affinity - the energy change associated with the addition of an electron • Can be an endothermic or exothermic process • Affinity tends to decrease as you go down • in a period Electrons farther from the nucleus experience less nuclear attraction Some irregularities due to repulsive forces in the relatively small p orbitals

  25. Electron Affinity - the energy change associated with the addition of an electron • Affinity tends to increase across a period • Affinity tends to decrease as you go down • in a period Electrons farther from the nucleus experience less nuclear attraction Some irregularities due to repulsive forces in the relatively small p orbitals

  26. Table of Electron Affinities

  27. Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons • Electronegativities tend to increase across • a period • Electronegativities tend to decrease down a • group or remain the same

  28. Electronegativity What is it? Electronegativity is the power of an atom to attract electron density in a covalent bond

  29. Electronegativity Pauling’s electronegativity scale The higher the value, the more electronegative the element Fluorine is the most electronegative element It has an electronegativity value of 4.0

  30. Electronegativity Pauling’s electronegativity scale

  31. Electronegativity

  32. Periodic Table of Electronegativities

  33. Ionic Radii Cations • Positively charged ions • Smaller than the corresponding • atom Anions • Negatively charged ions • Larger than the corresponding • atom

  34. Sizes of ions • Ions are atoms that have either gained or lost electrons (so that the # of electrons is not equal to the # of protons) • The size of an atom can change dramatically if it becomes an ion • E.g. when sodium loses its outer electron to become Na+ it becomes much smaller. Why? • Na+ is smaller than Na because it has lost its 3s electron. Its valence shell is now 2s22p6 (it has a smaller value of n) • Changing n values is one explanation for the size of ions. The other is …

  35. 9+ Sizes of ions: electron repulsion • When an atom becomes a –ve ion (adds an electron to its valence shell) the repulsion between valence electrons increases without changing ENC • Thus, F– is larger than F • Valence electrons push each other away

  36. Table of Ion Sizes

  37. electronegativity Summary of Periodic Trends electronegatvity Ionic radius Ionic radius

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