U3.1 Periodic Trends PDF

Periodic Trends We will explain observed trends in Atomic Radius size Ionic Radius lose e– Ionization energy Zeff & shielding (Effect of Protons) (Effect of Energy Levels) (explains A...

Periodic Trends We will explain observed trends in Atomic Radius size Ionic Radius lose e– Ionization energy Zeff & shielding (Effect of Protons) (Effect of Energy Levels) (explains ALL periodic trends and properties)  The Anatomy of an Atom Recall that: Oppositely charged particles attract Protons are positively towards one another while particles charged and located in with the same charge repel each other. the nucleus. Neutrons have no charge and are located in the nucleus. Electrons are negatively charged and are located outside the nucleus. Attractions and Repulsions of Subatomic Particles Coulomb’s law is used to calculate the force q1q2 (either attractive or repulsive) between two F= r2 charged particles. F : Force q1 : charge on particle 1 q2 : charge on particle 2 r : distance between charges Larger charges mean more electrostatic force between particles. Larger distance between particles means less electrostatic force between particles. Zeff & Shielding Zeff = Z − S effective nuclear charge, (Zeff): Z = nuclear charge (+proton’s) S = shielding (core e–’s) shielding shielding: Zeff inner core electrons shield outer e–’s attraction from nuclear attraction. Z = +11 Zeff = +1 Na atom Atomic Radius =shield, Zeff , attraction decreases across a period -due to increasing Zeff (more protons) increases down a group -due to increasing shielding (more energy levels) =Zeff shield attraction  Ionic Radius e– e– Na+ Cations are smaller than neutral atoms. outermost electron(s) are removed and loses a shell core shell closer to nucleus inner e–’s shielded (Zeff)  Ionic Radius e– e– e– Anions are larger than their parent atoms. electrons are added and repulsions are increased (=Z & =shielding) eff Arrange the following species by increasing size: Ar, K+, Ca2+, S2–, Cl– Ca2+ < K+ < Ar < Cl– < S2–  Ionization Energy (IE) energy required to remove an electron more energy to remove each electron IE1 < IE2 < IE3, … look for a huge jump in IE once all valence e–’s are removed, the next e– is on an inner level with attraction (shielding & Zeff). huge jump in IE4 b/c 4th e– on inner level (must have 3 valence e–’s)  Trends in First IE IE tends to… =shield increases across a period -due to increasing Zeff Zeff (more protons) attraction decreases down a group -due to shield increasing shielding =Zeff (more energy levels) attraction Group 1: Alkali Metals lowest IE’s (lose e–’s easily) Zeff more reactive down a group b/c…shielding causes att. & IE, easier to lose e– Group 2: Alkaline Earth Metals low IE’s, but not as low as alkali metals. less reactive than alkali metals (Zeff , att. & IE), but more reactive down the group. (shielding causes att. & IE) Group 17 or 7A: Halogens high IE’s (don’t lose e–’s easily) (Zeff , att.) large EN (attract e–) (Zeff , att.) more reactive at top of a group b/c… shielding causes att. & EN, easier for nonmetals to attract e– Group 18 or 8A: Noble Gases UNREACTIVE (mostly) b/c… Zeff , att. (no lose e–), HUGE IE’s b/c…… and filled valence shell (no gain e–) Monatomic gases