Chemistry Chapter 6

1. Chemistry Chapter 6 The Periodic Law

3. Mendeleev’s Periodic Table Dmitri Mendeleev

4. A Spiral Periodic Table “Mayan” Periodic Table Triangular Periodic Table

5. Chinese Periodic Table Giguere Periodic Table Stowe Periodic Table

6. Orbital filling table

7. Periodic Law • Chemical properties are functions of an elements atomic number

8. Periodic Table with Group Names

9. 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

10. Atomic Radius by period

11. Atomic Radius – Group He Be Mg Ca

12. Atomic Radius - Period The positive charge in the nucleus increases with each new proton added – increasing the nuclear charge. The negative charge on the outermost electron shell increases with each new electron added The built-up nuclear charge attracts the built-up electron charge making the outer shell pull in closer to the nucleus as protons & electrons are added.

13. Electron Shells & Size

14. Check your knowledge

15. Ions During chemical bonding atoms become stable by forming a noble gas electron configuration. Atoms will gain or lose electrons during bonding to achieve a noble gas electron configuration. If an atom accepts or donates (gains or loses) electrons the atom will become an ion. An ion is a charged particle. • Positive ions are called cations. • Negative ions are called anions.

16. Ionization of Magnesium Mg + 738 kJ  Mg+ + e- 1st ionization energy Mg+ + 1451 kJ  Mg2+ + e- 2nd ionization energy Mg2+ + 7733 kJ  Mg3+ + e- 3rd ionization energy

17. Ionization Energy • Ionization Energy is the measure of energy in KJ/mole that are required for removing an electron. • The first ionization energy is the energy needed to remove one electron from the outermost shell. • The second ionization energy is the energy needed to remove a 2nd electron from the outermost shell after the 1st one has been removed.

18. Removing An Electron makes the atom an ion

19. Ionization Energy – Trends

20. Predicting Ionic Charges Group 18: Stable Noble gases do not form ions!

21. Predicting Ionic Charges Group 1: Lose 1 electron to form 1+ ions H+ Li+ Na+ K+

22. Predicting Ionic Charges Group 2: Loses 2 electrons to form 2+ ions Be2+ Mg2+ Ca2+ Ba2+ Sr2+

23. Predicting Ionic Charges Group 13: Loses 3 electrons to form 3+ ions B3+ Al3+ Ga3+

24. Predicting Ionic Charges Group 14: Lose 4 electrons or gain 4 electrons? Neither! Group 13 elements rarely form ions.

25. Predicting Ionic Charges Nitride N3- Group 15: Gains 3 electrons to form 3- ions P3- Phosphide As3- Arsenide

26. Predicting Ionic Charges Oxide O2- Gains 2 electrons to form 2- ions Group 16: S2- Sulfide Se2- Selenide

27. Predicting Ionic Charges F1- Fluoride Br1- Bromide Group 17: Gains 1 electron to form 1- ions Cl1- Chloride I1- Iodide

28. Charge on Ions ? ? The charge on an ion is also called its Oxidation #

29. Ionic Radii Cations • Positively charged ions • Cations are smaller than their parent atoms because they lose electron(s) and also lose their outer energy level. Anions • Negatively charged ions • Anions are larger than their parent atoms because they keep their outer energy level and they cannot hold their gained electron(s) as tightly as their natural electrons

30. Unbalanced charges In normal atoms: the # of Protons = the # of electrons (positive and negative charges cancel each other out) When an atom loses or gains an electron during chemical bonding the charges are unbalanced. • If an electron is lost, there is one fewer electron • # ofprotons > # of electrons • there is more positive charge than negative charge The atom that donated the electron will now have one more proton than electrons is said to have a charge of plus 1 (+1).

31. Common MetalsCations

32. Common NonmetalsAnions

33. Atomic and Ionic radii Comparison

34. 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

35. Ionization Energies

36. 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 Think Greedy!!!!!

37. Electronegativity Electronegativity is the ability of an atom to attract an electron • Shielding decreases the ability to attract an electron. • Cations do not want to attract electrons because they want, instead, to lose electrons (the opposite thing). Therefore cations have a low ability to attract an electron. • Anions easily attract electrons because they want to gain electrons, making them high in electronegativity. • Noble gases, with their filled sublevels, do not want to attract electrons so they have no ability to attract electrons and therefore have no electronegativity.

38. Electronegativity

39. Shielding Effect The nuclear charge is shielded by the filled inner electron shells + + + + The nucleus does not have enough attraction to hold on to electrons far away from the nucleus.

40. Atomic Radius Slight Decrease Increase

41. Ionization Energy Increase Decrease

42. Electronegativity Increase Decrease