Periodic Trends

1. Periodic Trends

2. Periodic Trends • What is a trend? • A trend is the general direction in which something tends to move.

3. Periodic Trends • The elements on the Periodic Table of Elements show many trends in their physical and chemical properties. • Across the rows (periods or series) • Down the columns (groups or families)

4. Atomic Radius • ½ the distance between the nuclei of 2 like atoms in a diatomic molecule • The atoms of the 8 main groups are shown here.

5. Atomic Radius - Groups • Atomic radius increases as you move down a group • Why? • More electrons in more Principal Energy Levels • Atomic size increases

6. Atomic Radius - Periods • Atomic radius decreases as you move across a period • Why? • (-) electrons increase, but so do (+) protons !!! • Increased (+) nuclear charge pulls the (-) electrons closer to the nucleus • Atomic size decreases

7. Atomic Radius - Periods • The size trend in periods is less pronounced than in groups because of the electron shielding effect.

8. Shielding Effect • Reduction in effective nuclear charge on an electron that is caused by the repulsive forces of other electrons between it and the nucleus • In an atom with one electron, that electron experiences the full charge of the positive nucleus. However, in an atom with many electrons, the outer electrons are simultaneously attracted to the positive nucleus and repelled by the negatively charged electrons.

9. Atomic Radius - Graph

10. Atomic Radius • Click on source to see a short video • Source: http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/046_AtomicRadii.MOV

11. Atomic Radius

12. Ionic Radius • What are ions? • Ions are charged atoms, either + or - • Cations are positive ions • Cations form when atoms lose electrons • Anions are negative ions • Anions form when atoms gain electrons

13. Ionic Radius

14. Ionic Radius – Cations Group • Cations are smaller than their parent atoms. • Why? • By losing their valence electrons, they lose their entire valence shell • Cations are formed by the metals on the left side of the Periodic Table

15. Ionic Radius – Cations Groups • Ionic size increases as you move down a group for the same reason atomic size increases • Number of principal energy levels increases

16. Ionic Radius – Anions Groups • Anions are larger than their parent ions • Why? • When extra (-) electrons are added, extra (+) protons are NOT added to the nucleus • Effective nuclear attraction is less for the increases number of electrons

17. Ionic Radius – Anions Group • Ionic size increases as you move down a group for the same reason atomic size increases • Number of principal energy levels increases • Cations are formed by nonmetals on the right side of the Periodic Table

18. Ionic Radius - Periods • Just like their parent atoms… • Cations get smaller as you move from left to right • Anions get smaller as you move from left to right • Increased (+) nuclear charge pulls the (-) electrons closer to the nucleus

19. Ionic Radius

20. Ionization Energy • Energy is needed to remove an electron from an atom • The energy needed to overcome the attraction of the nuclear charge and remove an electron (from a gaseous atom) is called the Ionization Energy

21. 1st Ionization Energy • The energy needed to remove the 1st electron from an atom is the 1st Ionization Energy

22. Factors AffectingIonization Energy • Atomic Radius • Smaller atoms hang on to valence electrons more tightly, and so have higher ionization energy

23. Factors AffectingIonization Energy • Charge • The higher the positive charge becomes, the harder it is to pull away additional electrons • Second ionization energy is always higher than the first

24. Factors AffectingIonization Energy • Orbital Type • It's easier to remove electrons from p orbitals than from s orbitals, which are “deeper”

25. Factors AffectingIonization Energy • Electron Pairing • Within a subshell, paired electrons are easier to remove than unpaired ones • Reason: repulsion between electrons in the same orbital is higher than repulsion between electrons in different orbitals

26. Factors AffectingIonization Energy • Electron Pairing – Example • On the basis of gross periodic trends, one might expect O to have a higher ionization energy than N. However, the ionization energy of N is 1402 kJ/mol and the ionization energy of O is only 1314 kJ/mol. • Taking away an electron from O is much easier, because the O contains a paired electron in its valence shell which is repelled by its partner.

27. 1st Ionization Energy

28. 2nd, 3rd, 4th Ionization Energies, etc. • Subsequent electrons require more energy to remove than the first electron • How much more energy is needed depends on what energy levels and orbitals the electrons are in

29. 2nd, 3rd, 4th Ionization Energies, etc. • Source: http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch7/ie_ea.html

30. Ionization Energy • Click on source to see a short video. • Source: http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/047_IonizationEner.MOV

31. Electronegativity • Ability of an atom to attract electrons toward itself in a chemical bond

32. Electronegativity • The difference between the electronegativities of two atoms will determine what kind of bond they form • Linus Pauling used an element's ionization energy and electron affinity to predict how it will behave in a bond. • The more energy it takes to pull off the outer electron of an atom, the less likely it is to allow another atom to take those electrons. The more energy the atom releases when it gains an electron, the more likely it is to take electrons from another atom in bonding. These two energies were used to compute a numerical score.

33. Electronegativity - Periods • Electronegativity increases going left to right across the periodic table. • Fluorine's high nuclear charge coupled with its small size make it hold onto bonding electrons more tightly than any other element. Lithium has a lower nuclear charge and is actually larger than fluorine. Its valence electron is not tightly held and it tends to surrender it in chemical bonds.

34. Electronegativity - Groups • Electronegativity decreases going down a group • The bonding electrons are increasingly distant from the attraction of the nucleus

35. Electronegativity

36. Electronegativity

37. Electron Affinity • The electron affinity is a measure of the energy absorbed when an electron is added to a neutral atom to form a negative ion. • Most elements have a negative electron affinity. This means they do not require energy to gain an electron; instead, they release energy. • Atoms more attracted to extra electrons have a more negative electron affinity. • The more negative the value, the more stable the ion is.

38. Electron Affinity • Click on source to see a short video. • Source: http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/049_ElectrAffinity.MOV

39. Electron Affinity • Electron affinity is essentially the opposite of the ionization energy.

40. Electron Affinity - Trends • Click on source to see a short video. • Source: http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/050_PeriodElectron.MOV

41. Image Sources • http://www.Chem4kids.com • http://images.encarta.msn.com • http://antoine.frostburg.edu • http://www.webelements.com • http://cwx.prenhall.com • http://www.800mainstreet.com • http://intro.chem.okstate.edu/1215

42. Credits • PowerPoint: Adela J. Dziekanowski