IB SL Chemistry Unit 3

UNIT THREE ● Periodicity ○ Periodic table is separated into 5 groups ■ Group 1- Alkali Metals ● Lithium, Sodium, Potassium, Rubidium, Caesium, Francium ● High reactivity, will create explosion in water ■ Group 2- Alkaline Earth Metals ● Beryllium, Magnesium, Calcium, Strontium, Barium, Radium ● Reactive, but less than group 1 ■ Group 3-12 ● Transition metals ● Characterized by having electrons end in the D orbital ● From Scandium to Copernicium ● Not very reactive, mainly used in construction ■ Group 17- Halogens ● From Fluorine to Astatanine ● Most reactive nonmetals ■ Group 18- Noble Gases ● Most unreactive elements on the table due to its full valence shell ● Helium to Radon ○ Periodic blocks ■ Blocks are a way to characterize periodic table elements by what their valence shell is. ■ Transition metals are characterized by their D orbital valence electrons. ○ Atomic Radii ■ Atomic radius is distance from nucleus to valence shell. ■ Atomic radius decreases across a row and increases down a group. ■ This is because as you move across a row, both more protons and electrons are added. Protons will increase the nuclear charge of the atom and pull in the electrons tighter. ■ It increases down a group because a new group adds a new energy level of electrons. ○ Ionic radius ■ Metals will form cations, nonmetals will usually form anions. ■ Cations lose electrons, so the radius becomes stronger. Proton count stays the same, and pulls in less electrons. ■ Imagine a little girl walking 4 dobermanns. ■ If she loses a dobermann, she only has to control 3 of them. That means she can hold them closer to her and have a better grip on them. ■ Conversely, anions increase atomic radius. You add on electrons, so it takes the same amount of protons to pull in more electrons. The trend for ionic radius is that it will decrease across a row for the metals, jump up to a higher value once it hits a nonmetal, then decrease again. ■ Ionic radius will increase down a group, as nuclear charge will become stronger. ○ Nuclear charge- the amount of attraction between nucleus and valence shell ■ Will increase as you move up a group and move across a row. ○ First ionization energy ■ The energy needed to pull one electron off a gaseous atom. ■ Increases across a row, down a group. ● We observe peaks at all of the Noble Gases. This is because Noble gases have a full valence shell and are happy. Imagine trying to steal a candy bar from a kid. It will take a lot of effort. We see some exceptions in the row, mainly beryllium and boron, then oxygen and nitrogen. ● These can be proven by the electron configurations of each. ● Electronegativity ○ Defined as attraction of bonding for each element. ○ Francium has the lowest electronegativity, and it increases up a group and across a row until you get to Fluorine. ● Electron affinity ○ The energy released when one mole of electrons to a gaseous element. ○ Increases (becomes more exothermic as you go down a row) ○ Halogens have the highest electron affinity, they have room for one electron in their outer shell. ○ Decreases down a group ● Metallic character ○ Metallic character decreases across a row, and increases down a group ● Melting point ○ Generally is higher with metals, then decreases when it reaches nonmetals ● Oxide trends ○ Row three oxides decrease in metallic character as you move down the row. ○ Row three oxides go from basic to amphoteric to acidic. ○ Aluminum oxide is amphoteric, which means it can be both basic or acidic. ○ A common IB Question is to describe this change in character of oxides so please note this.