1. IIIIII Periodicity – the tendency to recur at regular intervals. For example: the return of the full moon every 28 days. Periodic Table & Trends

2. Periodic Law zWhen elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

3. Founding Scientists zRussian chemist Dmitri Mendeleev proposed 1 st periodic table where elements were arranged according to ATOMIC MASS. (1869) zEnglish chemist Henry Moseley revised periodic table and arranged elements according to ATOMIC NUMBER. (1913)

4. Chemical Reactivity zGroups or Families ySimilar valence e - within a group result in similar chemical properties

5. Chemical Reactivity zAlkali Metals zAlkaline Earth Metals zTransition Metals zHalogens zNoble Gases

6. zAtomic Radius ysize of atom © 1998 LOGAL zFirst Ionization Energy yEnergy required to remove one e - from a neutral atom. © 1998 LOGAL Periodic Trends

7. zIncreases to the LEFT and DOWN Atomic Radius

9. zDown a Group: Atomic Radius increases. yReason: Increasing the number of energy levels (shells) zFrom left to right across a Period: Atomic Radius decreases yReason: There is an increase in the number of protons in the nucleus causing there to be more of an effective nuclear charge, thus, a greater pull on the orbiting electrons in question. Atomic Radius

10. Effective Nuclear Charge zThe effective nuclear charge (often symbolized Z eff or as Z*) is the net positive charge experienced by an electron in a multi-electron atom. The term "effective" is used because the shielding effect of negatively charged electrons prevents higher orbital electrons from experiencing the full nuclear charge by the repelling effect of inner- layer electrons. The effective nuclear charge experienced by the outer shell electron is also called the core charge. It is possible to determine the strength of the nuclear charge by looking at the oxidation number of the atom.

11. Calculating the effective nuclear charge zIn an atom with one electron, that electron experiences the full charge of the positive nucleus. In this case, the effective nuclear charge can be calculated from Coulomb's law. zHowever, in an atom with many electrons the outer electrons are simultaneously attracted to the positive nucleus and repelled by the negatively charged electrons. The effective nuclear charge on such an electron is given by the following equation: Z eff = Z – S zZ is the number of protons in the nucleus (atomic number), and S is the average number of electrons between the nucleus and the electron in question (the number of non-valence electrons).

12. Electronegativity zElectronegativity – a measure of the ability of an atom in a bond to attract electrons. yImagine bonded atoms as being locked in a tug-of-war competition over the shared valence electrons. The atom’s ability to “tug” is its electronegativity. yAmerican chemist Linus Pauling designed a scale based on fluorine (F) and assigned it a value of 4.0 (this is the most electronegative element).

13. Electronegativity zFrom left to right across a Period: Electronegativity increases yReason: The number of protons in the nucleus increases, which increases the attractive force between the nucleus and the valence electrons zDown a Group: Electronegativity decreases zReason: Shielding – this is caused by an increase in the number of energy levels.

14. Electronegativity zElectronegativity yIncreases UP and to the RIGHT

15. zIonization Energy - the energy required to remove an electron from an atom. zFrom left to right across a Period: Ionization Energy increases. yReason: More protons in the nucleus creating a greater effective nuclear charge on orbiting electrons. zDown a Group: Ionization Energy decreases. yReason: Shielding Ionization Energy

16. zFirst Ionization Energy yIncreases UP and to the RIGHT Ionization Energy

17. Electron Affinity zThe Electron affinity of a molecule or atom is the energy change when an electron is added to the neutral atom to form an anion (negatively charges ion). This property can only be measured in an atom in its gaseous state.

18. zMelting/Boiling Point yHighest in the middle of a period. Melting/Boiling Point

19. zIonic Radius yCations (+) xlose e - xsmaller © 2002 Prentice-Hall, Inc. yAnions (–) xgain e - xlarger Ionic Radius

20. zWhich atom has the larger radius? yBe orBa yCa orBr Examples

21. zWhich atom has the higher 1st Ionization Energy? yNorBi yBa orNe Examples

22. zWhich atom has the higher melting/boiling point? yLiorC yCrorKr Examples

23. zWhich particle has the larger radius? ySorS 2- yAlorAl 3+ Examples