CHEM 101 Lecture Notes Week 4
CHEM 101 Lecture Notes Week 4 CHEM 101
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This 6 page Class Notes was uploaded by spencer.kociba on Friday October 14, 2016. The Class Notes belongs to CHEM 101 at Drexel University taught by Monica Ilies in Fall 2016. Since its upload, it has received 13 views.
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Date Created: 10/14/16
Spencer Kociba CHEM 101 Lecture notes Lecture 10/11/16 ○ Nuclear charge increases as the distance between the electron of interest and the nucleus decreases ● “D” has 5 pairs of electrons→ 10 electrons in each d level ● 1s 2s 2 p 3s 3p 2 ○ 1s through 2p are the core electrons, 3s and 3 p are the valence electrons (outermost shell) ● ***NOTE d orbitals are not included in valence electron count ○ Ex. ...4s 3d3 → the number of valence electrons=2 ○ (only s and p levels count for the valence electrons) ● D block= Rows 47, Columns 312 ● F block=Atomic numbers 58103 2 2 6 2 5 ● Cl=1s 2s 2 p 3s 3p 2 5 ○ Or [Ne]3s 3 p ○ ^^this is called Noble Gas Notation, it’s just shorthand when writing out the electron configurations for elements with many orbital levels ● Transition metals electron configuration ○ Expected 2 4 ■ Cr= [Ar]4s 3d ■ 2 9 Cu=[Ar]4s 3d ○ Actual ■ Cr=[Ar]4s 3d 1 5 ■ Cu=[Ar]4s 3d1 10 ○ Transition elements’ actual electrons and electron configuration are different from the expected and theoretical ■ Electrons try to fill/empty the outermost d shell to go into a lower energy state ■ Ex. Cu valence electron configuration= 4s 3d 10 ● D is filled even though s is not ● Halogens=column 17 (highly reactive because they only need one electron to complete their outer shell ● Alkali metals=column 1 (highly reactive because they only need to lose one electron) ● Alkaline earth metals=column 2 Spencer Kociba CHEM 101 Lecture notes Lecture 10/13/16 Concepts: trends in the periodic table, effective nuclear charges, covalent diameter/radius and Van der Waals diameter/radius ● Effective nuclear charge ○ Z❑ effS ■ Z❑ effis the effective nuclear charge ■ Z is the actual nuclear charge (protons) ■ S is the charge screened by other electrons (electrons) ■ → =increasing in this direction ● Trends in the periodic table ○ Ionization Energy (IE or E❑ ) i ■ Qualitatively defined as the amount of energy necessary to remove an electron of an isolated gaseous atom to form a cation ■ Group 2 has slightly higher IE values than expected because it is easier to remove a single electron from an orbital than it is to remove an electron by splitting up a pair ■ Group 16 has slightly lower IE values than expected ■ When you get a full or an empty shell, the IE increases dramatically because you are left with a noble gas core (where removing the core electrons takes A LOT of energy) ■ → =increasing in this direction ■ Remember: there are exceptions to this rule, but generally speaking it follows this trend ○ Electron Affinity ■ An element’s likelihood of gaining an electron ■ Trends ● Noble gases are nearly 0 because their outermost shell is full ● Group 15 is less than 14 and 16 because at group 15, the electrons must begin pairing in the p level ● Group 2 (alkaline earth metals) are lower because the s level is already full so the new electron must go to the p level ○ Atomic Radii ■ NOTE: Ionization energy is the greatest when AR is the smallest ■ → =increasing in this direction ○ Density ■ → =increasing in this direction ○ Metallic Character ■ → =increasing in this direction ● Cations ○ Positively charged particle ○ Formed through the loss of electrons (related to Ionization Energy) ○ Usually metals form cations ○ # of protons > # of electrons ○ Decreasing of the electrons=protons attract the remaining electrons closer to the nucleus (the atomic radius decreases) ○ MEMORY TIP: cation/metals=+ charge ● Anion ○ # of protons < # of electrons ○ Element gains electrons (related to Electron Affinity) ○ Usually nonmetals form anions ○ Increasing the electrons means the protons can’t pull in the electrons as well so the electrons spread out and the atomic radius increases ○ MEMORY TIP: prefix an= non/opposite=negative charge/nonmetals
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