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  2. 2. Key Concepts• Effective Nuclear Charge: The attractive force exerted by the atoms positively charged nucleus (+) on an electron (- ). • Holds the electron in orbit. • Opposites attract. • As you add protons (+) the Effective Nuclear Charge grows larger.• Electron Shielding: The reduction in the attractive force (Effective Nuclear Charge) between a positively charged nucleus and its outermost electrons due to the cancellation of some of the positive charge by the negative charges of the inner electrons. • Each primary energy level added moves the valence electrons farther from the nucleus. • Each successive energy level also puts more inner shell electrons “in the way” between the nucleus and valence (outer) shell. • Both of which decrease the amount of attractive force the nucleus can exert.
  4. 4. Atomic Radius• The exact size of an atom is hard to determine.• The volume the electrons occupy is thought of as an electron cloud, with no clear-cut edge.• One method for calculating the size of an atom involves calculating the bond radius, which is half the distance from center to center of two like atoms that are bonded together.
  5. 5. Atomic Radius Increases as You Move Down a Group•As you proceed from one element down to the next in a group, a new principal energy level is added.•The addition of another level of electrons increases the size, or atomic radius, of an atom.•Effect is further magnified by electron shielding.
  6. 6. Atomic Radius Decreases as You Move Across a Period • As you move from left to right across a period, each atom has one more proton and one more electron than the atom before it has. • All additional electrons go into the same principal energy level—no electrons are being added to the inner levels – changes in electron shielding are minimal. • As the nuclear charge increases across a period, the effective nuclear charge acting on the outer electrons also increases.  Pulls the electrons in tighter, creating a smaller atom.
  7. 7. Atomic Radii
  9. 9. Ionization• We’ve previously looked only at neutral atoms with equal numbers of protons and electrons.• Ionization: The creation of an atom with a net charge, by the removal or addition of electrons. – Ions have more or less electrons than neutral atoms.• Anion: An atom that has gained an electron and taken on a net negative charge.• Cation: An atom that has lost an electron and taken on a net positive charge.
  10. 10. Ionization Energy•The ionization energy is the energy required to remove an electron from an atom or ion. •The Higher the Ionization Energy the harder it is to remove an electron. A + ionization energy ® A + + e - neutral atom ion electron First ionization energy: +1 Second ionization energy : +2
  11. 11. Multiple Ionization Energies• Multiple ionization energies: If you want to pull off more than one electron from an atom, more energy is required for each additional one you want to grab.• Once you’ve reached a noble gas configuration (for example, once magnesium has lost two electrons to become like neon), any further electrons you pull off will require a huge amount of energy. First ionization energy: +1 Second ionization energy : +2
  12. 12. Ionization Energy Decreases as You Move Down a Group• Each element has more occupied energy levels than the one above it has. • The outermost electrons are farthest from the nucleus in elements near the bottom of a group.• As you move down a group, each successive element contains more electrons in the energy levels between the nucleus and the outermost electrons- more electron shielding.• Therefore it gets easier to remove electrons as you go down.
  13. 13. Ionization Energy Increases as You Move Across a Period• Ionization energy tends to increase as you move from left to right across a period.• From one element to the next in a period, the number of protons and the number of electrons increase by one each- increases the effective nuclear charge.  A higher nuclear charge more strongly attracts the outer electrons in the same energy level, but the electron-shielding effect from inner-level electrons remains the same.
  14. 14. Electron Affinity• Basically the opposite of ionization energy.• The amount of energy released or absorbed when an atom accepts an electron giving it a negative charge. For most elements, energy is released when an atom adds an electron. This is also the measure of an element to attract an electron to form a negative ion.• How much it wants to gain an electron.• Electron affinity increases from left to right and decreases from top to bottom in a group or family.
  16. 16. Ion Radii• Negative Ions: Always larger than the neutral atom. Gaining electrons.• Positive Ions: Always smaller that the neutral atom. Loss of outer shell electrons.
  18. 18. Electronegativity• Electronegativity: is a measure of an atom’s attraction for another atom’s electrons when in a chemical compound. – It is an arbitrary scale that ranges from 0 to 4. • The units of electronegativity are Paulings. – The atom with the higher electronegativity will pull on the electrons more strongly than the other atom will. – Generally, metals are electron givers and have low electronegativities. – Nonmetals are electron takers and have high electronegativities. – What about the noble gases? • They have no electronegativity at all because they are noble gases, and don’t need any more electrons
  19. 19. Electronegativities of Some Elements Element Pauling scale F 4.0 Cl 3.0 O 3.5 N 3.0 S 2.5 C 2.5 H 2.1 Na 0.9 Cs 0.7
  20. 20. Electronegativity Decreases as You Move Down a Group• Electronegativity values generally decrease as you move down a group. • As you move from higher to lower in a group, the electronegativity decreases due to the increase in separation and effects of electron shielding. • Generally, if an atom doesn’t hold the electrons it already has very strongly (low ionization energy), it won’t want to grab electrons from other atoms.
  21. 21. Electronegativity Increases as You Move Across a Period• Electronegativity usually increases as you move left to right across a period.• As you proceed across a period, each atom has one more proton and one more electron—in the same principal energy level—than the atom before it has.  Therefore the effective nuclear charge increases across a period, resulting in an increase in electronegativity.• The increase in electronegativity across a period is much more dramatic than the decrease in electronegativity down a group.
  22. 22. Nuclear charge increasesShielding increasesAtomic radius increasesIonic size increasesIonization energy decreasesElectronegativity decreases Shielding is constant Summary Atomic Radius decreases Nuclear charge increases Electronegativity increases Ionization energy increases