The Periodic Table

Dimitri Mendeleev dev. the modern periodic system Placed elements in order of increasing atomic mass Reversed the order of some elements to keep similar elements in the same column Ex. Ar and K Predicted the existence and properties of elements that had not yet been discovered-this demonstrated the value of the table

Modern Periodic Table Elements are placed in order of increasing atomic number Periodic Law: - the phys. and chem. properties of the elements are periodic functions of their atomic numbers.

Overview of Periodic Table Horizontal rows are called periods Elements in a period have similar electron configurations. Ex. Li -1s22s1 B- 1s22s22p1 **the 2nd energy level is the outermost for both Vertical columns are called groups-elements in a group have similar properties Ex. Group 1 elements are all highly reactive

The table is divided into two distinct regions: metals and nonmetals Metals are in green, nonmetals in yellow, and metalloids in blue.

Properties of Metals Good conductors of heat and electricity Most are solids at room temp.- the only exception is Hg Malleable-can be bent Ductile- can be drawn into a wire

Properties of Nonmetals Poor conductors of heat and electricity Not malleable or ductile Most are gases or liquids at room temp.

Properties of Metalloids Have properties of both metals and nonmetals Often used as semiconductors

Main Group Elements Groups 1, 2, and 13-18 Sometimes called the representative elements.

Transition Elements Located in groups 3-12

Alkali Metals React w/water to produce alkaline solutions and have metallic properties Soft Have shiny surfaces Highly reactive Become plasmas in gaseous state All located in Group 1

Alkaline Earth Metals Harder, denser, and stronger than alkali metals Less reactive than alkali metals All located in Group 2

Halogens Halogen is derived from a Greek word that means “salt former” These elements combine w/metals to form salts Most reactive group of nonmetals All located in Group 17 Chlorine Bromine

Noble Gases Formerly known as inert gases because they were believed to be completely unreactive Are unreactive in nature because of their full outermost energy levels All located in Group 18

Lanthanides and Actinides Collectively known as the Rare Earth Metals Lanthanides have atomic #’s 58-71-are shiny, reactive metals Actinides have atomic #’s 90-103 and are all radioactive

Hydrogen The most common element in the universe Unique behavior due to the fact that it only has 1 proton and 1 electron Can react w/many elements

Periodic Trends

Many properties of elements change in predictable ways: There are six periodic trends we will talk about. Atomic Radii Ionic Radii Ionization Energy Electron Affinity Electronegativity Valence Electrons

Atomic Radius The distance from the center of an atom’s nucleus to its outermost electron Atoms get larger down a group Atoms get smaller across a period Why? As you move down a group, there are more principle energy levels. As you move across a period, the attraction between the increasing # of protons and electrons causes the atom to shrink.

Note: You have to ignore the noble gas at the end of each period. Because neon and argon don't form bonds, you can only measure their van der Waals radius - a case where the atom is pretty well "unsquashed"

Examples Who has the larger radius: Li or Rb Ca or Br Be or Ba Al or Ga Answers: 1. Rb 2. Ca 3. Ba 4. Ga

Ionic Radii One-half the diameter of an ion in an ionic compound Cation- positive ion Anion-negative ion The formation of a cation results in a loss of one or more electrons and leads to a decrease in radius. The formation of an anion results in a gain of one or more electrons and an increase in radius.

Examples: Which has the larger radius? Ca or Ca+2 Br or Br-1 Fe+2 or Fe+3 Answers: Ca Br-1 Fe+2

Ionization Energy The energy required to REMOVE AN ELECTRON. Why?? The increased pull of additional protons as you go across a period results in a stronger pull on electrons which means more energy to remove them. Down a group, the outer electrons are further from the pull of the protons in the nucleus which means less energy to remove them. The energy required to REMOVE AN ELECTRON. Increases across a period Decreases down a group

The ionization energy for the removal for a second or third electron will be higher. Examples: Who has the larger ionization energy? Ca or Kr Be or Ba Fe+2 or Fe+3 Answers: Kr Be Fe+3

Electron Affinity The energy change that occurs when an atom gains an extra electron. Becomes more negative across a period ( the more negative, the more likely an electron will be gained) Becomes less negative down a group ( meaning less likely to gain an electron).

Examples: Answers: Cl F Which has the more negative electron affinity and therefore is more likely to gain an electron? Mg or Cl F or At

Electronegativity Ability to attract electrons. Fluorine is the most electronegative element. Increases as you go across a period. Decreases down a group. Which has the higher electronegativity meaning it is more likely to gain an electron? F or Ne Cl or I Answers: F ( the trend stops at Noble Gases) Cl

H: 2. 1 Li: 1. 0 Be: 1. 5 B: 2. 0 C: 2. 5 N: 3. 0 O:3. 5 F: 4. 0 Na: 0 H: 2.1 Li: 1.0 Be: 1.5  B: 2.0 C: 2.5 N: 3.0 O:3.5 F: 4.0 Na: 0.9 Mg: 1.2 Al: 1.5 S: 1.8 P: 2.1 S: 2.5 Cl: 3.0

Valence Electrons Group # Valence 1 1 2 2 13 3 14 4 1 1 2 2 13 3 14 4 15 5 16 6 17 7 18 8 The electrons available to be lost, gained, or shared in the formation of a chemical bond-the outermost electrons.