1. Periodic Table & Periodicity Ms Piela Durfee High

2. Periodic Trends/ Periodicity  A periodic trend is a pattern observed on the periodic table for an atomic property  Each of the four trends have explanations for their group trend and their period trend

3. The 4 Main Periodic Trends are: Atomic RadiusIonization EnergyElectronegativityElectron Affinity

4. The Period Trend Explanation  When comparing elements in the same period, compare the effective nuclear charges (Z eff )  Effective nuclear charge is the net positive charge experienced by electrons in an atom

5. The Period Trend Explanation  The atoms on the right of the periodic table have higher effective nuclear charges (Z eff ) when compared to elements on the left  This is due to electrons being added to the same energy level. They are approximately the same distance away from the nucleus  In general, the further atoms are away from the nucleus, the less attracted they become

6. The Group Trend Explanation  When comparing atoms in the same group, compare the amount of electron shielding occuring  Electron shielding is where core electrons shield outer electrons from the charge of the nucleus  Thus, outer electrons are held less tightly because of electron/electron repulsion

7. The Group Trend Explanation  Atoms on the top of the periodic table have less electron shielding than atoms at the bottom  As you increase in the number of energy levels, more electron shielding occurs  This does NOT occur across a period as energy levels will not change

8. Atomic Radius  Atomic Radius is a measure of the size of the atom  Measured by the distance from the nucleus to the outermost electrons

9. Atomic Radius  Atomic Radii decreases moving across a period, and increases going down a group  For the period trend: with effective nuclear charge, the increased positive charge pulls electrons closer, causing the size to decrease  With the group trend, the increasing energy levels provide more electrons, which increase the size of the atom (electron shielding doesn’t really work)

10. Ionization Energy  Ionization energy is the energy required to remove an electron from an atom  Amount of energy increases as the number of ionizations occur (i.e. first ionization takes less energy than the second, and so on)

11. Ionization Energy  Ionization energy increases going across a period and decrease going down a group  With increasing effective nuclear charge, electrons are held more tightly, thus atoms on the right require more energy to remove an electron  With increasing electron shielding, electrons are held less tightly and thus decrease in IE

12. Graph of IE Periodic Trend

13. Electronegativity  Electronegativity is the ability of an atom in a molecule to attract shared electrons to itself  Think of electronegativity as a “tug of war”

14. Electronegativity  Electronegativity increases going across a period, and decreases going down a group  Due to increasing effective nuclear charge, atoms on the right hold electrons more tightly, causing them to have high EN  Due to electron shielding, atoms on thebottom tend to hold electrons more loosely, making them have low EN

15. Electronegativity  The noble gases are excluded from this trend as they tend not to bond with other atoms  This makes fluorine the most electronegative atom

16. Electron Affinity  Electron Affinity is the energy associated with the addition of an electron to an atom  The more negative the quantity, the more energy is released upon the addition of an electron

17. Electron Affinity  Electron affinity increases across a period and decreases going down a group  Due to increasing effective nuclear charge, atoms on the right tend to want to attract negative electrons more  Due to electron shielding, atoms on the bottom tend to hold electrons more loosely, making them have low EA