The Periodic Table and Periodic Law

Development of the Modern Periodic Table

 Until the 1790s, only 23 different elements were known, such as silver, gold, carbon, oxygen, etc  These elements had been known for hundreds or thousands of years  The 1800s brought forth many changes to the scientific community, including an explosion in the number of elements known

 By 1870, there were about 70 known elements  With the increase in the number of elements, came an increased need to organize them  The man given the most credit for organizing the elements is Dmitri Mendeleev  He organized the elements in order of increasing mass, and when he did this, he noticed a pattern in the properties of the elements

 Mendeleev arranged his periodic table similar to a winning configuration of solitaire  Within rows, the elements were arranged by increasing mass  Within columns, elements were arranged by similar chemical properties  The usefulness of Mendeleev’s table was confirmed by the discovery of new elements that matched predicted properties according to their location on the periodic table

 Mendeleev’s table was a good step forward, but it was not completely correct  It soon became apparent, based on certain chemical properties, that some elements were not in the correct order  It was not until 1913 when Henry Moseley discovered that each element has a unique positive charge that this was corrected

 Mendeleev did not know about subatomic particles when he made his periodic table  The fact that each element has a unique charge, while the mass can vary because of isotopes caused the periodic table to be rearranged  The modern periodic table would then now be arranged by the atomic number

 The modern periodic table is arranged by atomic number  When the periodic table is arranged in this fashion, there is a periodic repetition of chemical properties from row to row. This is called periodic law

 The rows on the periodic table are called periods  The columns on the periodic table are called groups or families  Groups can be numbered in two separate methods – groups can be designated with a number and a letter, or just numbered 1-18

 The A groups (IA-VIIIA) are called representative elements because they represent a wide variety of chemical and physical properties  The B groups (IB-VIIIB) are the transition elements. They transition from very metallic to less metallic

 Starting below boron, draw imagine a staircase down to the bottom of the periodic table  Elements below and to the left of this staircase, but not touching (aluminum is the exception) are metals  Metals are generally shiny, solid at room temperature, good conductors of heat and electricity, malleable, and ductile  The far left group (IA) is called the alkali metals  The next group (IIA) is called the alkaline earth metals

 Alkali metals and alkaline earth metals tend to be chemically active  The alkali metals are more reactive than alkaline earth metals  Within the group of alkali and alkaline earth metals, reactivity increases as we go down the group

 The B elements (transition elements) are divided into two groups: the transition metals and the inner transition metals  The inner transition metals are the two rows located below the main body of the periodic table  Inner transition metals (specifically the lanthanide series) are used as phosphors, substances that emit light when struck by electrons

 To the right and above (but not touching) the staircase are the nonmetals  Nonmetals are generally gases at room temperature, brittle, dull, poor conductors of heat and electricity, and nonductile  The group VIIA nonmetals are called the halogens – from the Greek halos meaning salt and genesis meaning formation  The halogens are extremely reactive and react readily with metals. Within the halogen group, fluorine is the most reactive and reactivity decreases as we move down the group  Group VIIIA are called the noble gases, which are extremely UNreactive

 The elements that border the staircase are the metalloids (only element touching the staircase that is not a metalloid is aluminum)  Metalloids have properties that can be similar to both metals and nonmetals, depending on temperature, other compounds present, etc.  Silicon is a good example. In a computer, silicon is used to conduct electricity in circuit boards. In baking, silicon is used to make heat resistant oven mitts

Classification of the Elements

 What is a valence electron?  Atoms in the same group have the same number of valence electrons  Because elements in the same group have the same number of valence electrons, they have similar chemical properties  Within a period, the number of valence electrons increases from left to right  The number of valence electrons in representative elements can be found on the periodic table. The number paired with the A heading indicates the number of valence electrons

 Remember electron configurations? Write the electron configuration for iron

 The electron dot diagram explains why we start a new row after each noble gas  Their outer electron level gets full, so anything added after needs to be added to a new energy level  The rows on the periodic table indicate what energy level those electrons are occupying

 The periodic table is divided into blocks, the s-, p-, d-, and f-block  These blocks are the same as the s, p, d, and f energy sublevels we talked about earlier

Periodic Trends

 Technically defined, atomic radius is the area where there is a 90% probability of finding electrons  Essentially, it is the size of the atom  Atomic radius increases from top to bottom within a group. Why?  Within a period, atomic radius decreases. Why?  It decreases because we are adding protons and electrons, but staying on the same energy level. This increases effective nuclear charge which pulls the electrons in tighter, making the atom smaller

 Atoms can gain or lose electrons  When they do this they form an ion, which is an atom that has a net positive or negative charge  When an atom loses electrons, it forms a positive charge and the radius decreases from its neutral atom  When an atom gains electrons, it forms a negative charge and the radius increases from its neutral atom  Within a period ionic radius of positive ions decreases, then increases as it changes from positive to negative ions, then decreases again  Within a group, ionic radius increases

 Ionization energy – the energy required to remove an electron from an atom  Ionization energy increases from left to right across a row. Why?  Ionization energy decreases from top to bottom. Why?  Atoms can have second or third ionization energies as well

 Electronegativity is the desire for electrons  Electronegativity increases from left to right across a period  Electronegativity decreases from top to bottom  The noble gases have very minimal electronegativity  The most electronegative element is fluorine, while the least is francium