Periodic Table and Periodic Trends

1. Periodic Table andPeriodic Trends

2. Periodic Table of the Elements • is an arrangement of the elements according to their properties. • It enables chemists to classify the elements so that it is possible to identify patterns and trends in their properties. • Many scientists have made significant contributions to the development of the modern Periodic Table,  

3. Johann Doberiener (1829) • Arranged elements in groups of three (triads) • For example: Cl (mass 35) Br (mass 80) I (mass 127) Note: Average mass of the extreme elements =mass of the middle element

4. Lothar Meyer (1867) • Father of the Modern Periodic table. • Based on increasing atomic mass • Chemical properties of elements reoccur in a periodic fashion.

5. Dmitri Mendelyeev (1869)(Russian chemist) • Like Meyer prepared a table based on increasing atomic mass. • Consisted of 17 columns with 63 elements • Widely accepted • Constantly revised as new masses were discovered • Left vacant spots for new elements (Ga, Cd,Ge,Tc, Re, Po) • Certain discrepancies as new masses were found: Ar and K, Co and Ni, Te and I

6. Henry Mosley (1913) • English scientist used x-rays to determine atomic number (number of protons) • Discovered that the properties of elements were periodic functions of their atomic numbers, • Many of the discrepancies disappeared.

7. Present arrangement • Based on increasing atomic number • Periodic Law: properties of the elements depend on the structure of the atom and vary with the atomic number in a systematic way • Horizontal rows are called periods (#1-7) Indicates the number of principal energy levels. Note: Lanthanide series part of period 6—long period and Actinide series part of period 7—unfinished period Properties of elements change within a period • Vertical columns are called groups or families (#1-18) Indicates the number of valence electrons. Elements within a group have similar properties. Ex.Na and Li

8. Element Key 0

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

10. 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)

11. 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.

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

13. 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

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

15. 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.

16. Atomic Radius - Graph

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

18. Atomic Radius

19. 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

20. Ionic Radius

21. 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

22. 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

23. 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

24. 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

25. 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

26. Ionic Radius

27. 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

28. 1st Ionization Energy • The energy needed to remove the 1st electron from an atom is the 1st Ionization Energy

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

30. 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

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

32. 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

33. 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.

34. 1st Ionization Energy

35. 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

36. 2nd, 3rd, 4th Ionization Energies, etc. • Source: http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch7/ie_ea.html

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

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

39. 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.

40. 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.

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

42. Electronegativity

43. Electronegativity

44. 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.

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

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

47. 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

48. 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

49. Credits • PowerPoint: Adela J. Dziekanowski