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Core Electrons

Core Electrons. These are the inner electrons of an atom . They are not exposed very much to the electrons of other atoms when chemical bonds are formed. Core Electrons.

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Core Electrons

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  1. Core Electrons These are the inner electrons of an atom. They are not exposed very much to the electrons of other atoms when chemical bonds are formed.

  2. Core Electrons These are the inner electrons of an atom. They are not exposed very much to the electrons of other atoms when chemical bonds are formed. Examples: Sc has the configuration [Ar]4s23d1

  3. Core Electrons These are the inner electrons of an atom. They are not exposed very much to the electrons of other atoms when chemical bonds are formed. Examples: Sc has the configuration [Ar]4s23d1 The core is1s22s22p63s23p6 which is represented by [Ar]

  4. The electron configuration of N is 1s22s22px2py2pz

  5. The electron configuration of N is 1s22s22px2py2pz This can be represented as [He]2s22px2py2pz

  6. The electron configuration of N is 1s22s22px2py2pz This can be represented as [He]2s22px2py2pz The core is 1s2 in this example.

  7. Valence Electrons

  8. Valence Electrons The outer shell is called the valence shell, and the electrons in it are called valence electrons.

  9. Valence Electrons The outer shell is called the valence shell, and the electrons in it are called valence electrons. Examples: K has the electronic configuration [Ar]4s1. The valence electron configuration is 4s1.

  10. Valence Electrons The outer shell is called the valence shell, and the electrons in it are called valence electrons. Examples: K has the electronic configuration [Ar]4s1. The valence electron configuration is 4s1. Al has the configuration [Ne]3s23p1. The valence electron configuration is 3s23p1.

  11. The valence electrons are the most important. They are the electrons that most strongly influence the nature and formation of chemical bonds.

  12. Atomic Radii and Ionic Radii

  13. Atomic Radii and Ionic Radii One way to estimate the size of an atom is to measure the atomic radius, defined to be half the distance between two identical atoms in a molecule.

  14. Atomic Radii and Ionic Radii One way to estimate the size of an atom is to measure the atomic radius, defined to be half the distance between two identical atoms in a molecule. For example, the distance between two identical atoms (measured from the nuclei) in an I2 molecule is 2.66 Å. The radius of each iodine atom is taken to be 1.33 Å.

  15. The atomic radius decreases as we move across a period(there are some minor exceptions to this). The increase in the nuclear charge as we move across a period results in a shrinkage of the atomic radius – that is, the charge cloud is more strongly attracted to the nucleus.

  16. As we go down the periodic table, the atomic radius increases with increasing atomic number. The orbital size increases with increasing principal quantum number n.

  17. As we go down the periodic table, the atomic radius increases with increasing atomic number. The orbital size increases with increasing principal quantum number n. The corresponding increase in the nuclear charge does not decide the issue of size as we go down the periodic table.

  18. As we go down the periodic table, the atomic radius increases with increasing atomic number. The orbital size increases with increasing principal quantum number n. The corresponding increase in the nuclear charge does not decide the issue of size as we go down the periodic table. The size of atoms plays an important role in the nature of chemical bonds.

  19. Ionic Radius

  20. Ionic Radius Ionic radius: The radius of a cation or an anion as measured in an ionic compound.

  21. Ionic Radius Ionic radius: The radius of a cation or an anion as measured in an ionic compound. Anions (single atom ones) have larger radii than cations in the same period.

  22. Consider the isoelectronic species: Na+, Mg2+, Al3+

  23. Consider the isoelectronic species: Na+, Mg2+, Al3+ The smallest radii occurs for Al3+. The next smallest is Mg2+, followed by Na+.

  24. Consider the isoelectronic species: Na+, Mg2+, Al3+ The smallest radii occurs for Al3+. The next smallest is Mg2+, followed by Na+. In the trivalent cation, the electron density is pulled inward towards the nucleus most strongly by the +3 charge on the nucleus.

  25. For the anions O2- and F-, the oxide ion is the larger, because the extra electrostatic repulsion between the electrons in O2- will spread out the electron density to a greater extent than in F-.

  26. For the anions O2- and F-, the oxide ion is the larger, because the extra electrostatic repulsion between the electrons in O2- will spread out the electron density to a greater extent than in F-. The N3- ion is larger than the oxide ion.

  27. The Periodic Table

  28. The Periodic Table Some trends

  29. Definition: The ionization energy is the energy required to remove an electron from one mole of a substance in its ground state in the gas phase. For example, for substance X, X(g) X(g)+ + e- This defines the first ionization energy.

  30. The smaller the ionization energy, the easier a cation may be formed.

  31. The smaller the ionization energy, theeasier a cation may be formed. The ionization energy gives information about the chemical reactivity of an element.

  32. Element First ionization energy (kJ/mol) Li 520 increasing reactivity

  33. Element First ionization energy (kJ/mol) Li 520 Na 496 increasing reactivity

  34. Element First ionization energy (kJ/mol) Li 520 Na 496 K 419 increasing reactivity

  35. Element First ionization energy (kJ/mol) Li 520 Na 496 K 419 Rb 403 increasing reactivity

  36. Element First ionization energy (kJ/mol) Li 520 Na 496 K 419 Rb 403 Cs 376 increasing reactivity

  37. Element 1st IE Li 520 Na 496 K 419 (units are kJ/mol)

  38. Element 1st IE 2nd IE Li 520 7300 Na 496 4560 K 419 3052 (units are kJ/mol)

  39. Element 1st IE 2nd IE 3rd IE Li 520 7300 11808 Na 496 4560 6900 K 419 3052 4410 (units are kJ/mol)

  40. Element 1st IE 2nd IE 3rd IE Li 520 7300 11808 Na 496 4560 6900 K 419 3052 4410 (units are kJ/mol) From this table it is clear why we do not have Li2+, Na2+, K2+, or Li3+, Na3+, K3+ as common cations.

  41. Element 1st IE(kJ/mol) He 2370 Ne 2080 Ar 1520 Kr 1350 Xe 1170

  42. Element 1st IE Be 899 Mg 738 Ca 590 (units are kJ/mol)

  43. Element 1st IE 2nd IE Be 899 1757 Mg 738 1450 Ca 590 1145 (units are kJ/mol)

  44. Element 1st IE 2nd IE 3rd IE Be 899 1757 14850 Mg 738 1450 7730 Ca 590 1145 4900 (units are kJ/mol)

  45. Element 1st IE 2nd IE 3rd IE Be 899 1757 14850 Mg 738 1450 7730 Ca 590 1145 4900 (units are kJ/mol) From this table it is clear why we do not have Be3+, Mg3+, Ca3+ as common cations.

  46. Electron Affinity

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