The Periodic Table

u Horizontal rows are called periods u There are 7 periods

u Vertical columns are called groups. u Elements are placed in columns by similar properties. u Also called families

main group elements – s-block and p-block

Group 1 are the alkali metals Group 2 are the alkaline earth metals 1 2

Group 13 Boron Family Group 14 Carbon Family 1314

Group 15 Nitrogen Family Group 16 Oxygen Family 1516

u Group 17 is called the Halogens u Group 18 are the noble gases 17 18

The d-block - groups 3-12 are called the transition elements

The f-block consists of the Inner Transition Elements are the Lanthanides and Actinides

The Periodic Table Booklet

Page 1- History u Cannizzaro u Dmitri Mendeleev u Moseley

s2s2 s1s1 S- block u Page 2 – Introduction to the s-block (pp ) u Page 3 – Group 1 Alkali Metals Characteristics, group diagram Uses and important facts (pp ) u Page 4 – Group 2 Alkaline Earth Metals Characteristics, group diagram Uses and important facts (pp )

P-block u Page 5 – Introduction to the p-block (pp ) u Pages 6-10 – Groups Characteristics, group diagram Uses and important facts (pp )

Transition Metals -d block Page 11 – Introduction to the d-block (pp ) Pages – Groups 3-12 Characteristics Uses and important facts (pp )

(1)Atomic Size u Atomic Radius = half the distance between two nuclei of a diatomic molecule. } Radius

Atomic Size Group trends  Atomic radii increases as we go down a group u Why -Each atom has another energy level H Li Na K Rb

Atomic Size Period Trend u the atomic radii decrease as you go across a period u Why -More positive nuclear charge holds energy levels tighter NaMgAlSiPSClAr

Ionization Energy (IE) u The amount of energy required to completely remove an electron from a gaseous atom.

IE Group trends u IE decreases as you go down a group u Why -The electron is further away in a higher energy level and easier to remove

IE Period Trend u IE generally increases from left to right across a period u Why - Increasing + Nuclear Charge holds electrons tighter u ***Important note: u Metals have a lower ionization than nonmetals – This is the reason for the high reactivity of Gr. 1 metals. u Francium is the most reactive metal

Electron Affinity u The energy change associated with adding an electron to a gaseous atom.

EA Group Trend u Decrease as we go down a group. u Why -atomic radius is bigger and the energy level where the electron is added is farther away from the + nucleus

EA Period Trend u Increase from left to right across a period u Why – you add less electrons u *** Important note: u This is the reason for the high reactivity of Gr 17 elements u Fluorine is the most reactive of the nonmetals

Electronegativity u The ability of an atom in a chemical compound to attract an electron u High electronegativity means it pulls the electron toward it.

Electronegativity Trend u Same trends and reasons as EA u Decreases or stays the same down a group u Increase from left to right across a period

Metallic Properties u Decrease from left to right across a period u Increase as we go down a group. u Most metallic elements are at the bottom left u Most nonmetallic elements are at the top right

Ionization energy, electronegativity Electron affinity, DECREASEDECREASE INCREASE

Atomic size Metallic properties DECREASES INCREASESINCREASES

The ionization energy is lower for which element? a.sodium, cesium b. potassium, calcium Which is more chemically active? a.Mg, Cab. Cl, Br Which element is the biggest in size (radius)? a.K or Frb. K or Cac. F or At Which element can attract electrons easier? a. Mg, S b. N, F c. O, Ted. Cl, I

Ionic Size u Cations form by losing electrons. u Cations are smaller that the atom they come from. u Metals form cations. u Cations of representative elements have noble gas configuration.

Ionic size u Anions form by gaining electrons. u Anions are bigger that the atom they come from. u Nonmetals form anions. u Anions of representative elements have noble gas configuration.

Configuration of Ions u Ions always have noble gas configuration. u Na is 1s 2 2s 2 2p 6 3s 1 u Forms a +1 ion - 1s 2 2s 2 2p 6 u Same configuration as neon. u Metals form ions with the configuration of the noble gas before them - they lose electrons.

Configuration of Ions u Non-metals form ions by gaining electrons to achieve noble gas configuration. u They end up with the configuration of the noble gas after them.

Group trends u Adding energy level u Ions get bigger as you go down. Li +1 Na +1 K +1 Rb +1 Cs +1

Periodic Trends u Across the period nuclear charge increases so they get smaller. u Energy level changes between anions and cations. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1

Size of Isoelectronic ions u Iso - same u Iso electronic ions have the same # of electrons u Al +3 Mg +2 Na +1 Ne F -1 O -2 and N -3 u all have 10 electrons u all have the configuration 1s 1 2s 2 2p 6

Size of Isoelectronic ions u Positvie ions have more protons so they are smaller. Al +3 Mg +2 Na +1 Ne F -1 O -2 N -3