Chemistry Lab/Periodicity

Periodicity was a topic for the event Chemistry Lab in 2012 and 2013.

Overview
Periodicity, the repetition of patterns at regular intervals, is the most important property of the Periodic Table. In the Periodic Table, pure elements are ordered in increasing order of atomic number, which is the number of protons in nucleus. Other important factors in determining the arrangement of the elements include electron configurations and recurring chemical and physical properties. These periodic trends may increase, decrease, or stay consistent along a row or column. The periodic table also contains four rectangular blocks with approximately similar properties. Metals tend to be found more towards the left of the table while nonmetals tend to be found on the right, with metalloids in between, following a "staircase" pattern.

History
The first widely-accepted arrangement of the elements was created by the Russian chemistry professor Dimitri Mendeleev in 1869. Previous efforts to order the elements had been attempted, but none had proven to be very successful. Mendeleev's table grouped the elements into eight columns, rather than the now-conventional 18 column arrangement, in increasing order of atomic mass instead of atomic number. Elements with similar chemical properties were grouped into families, while Mendeleev boldly left spaces blank for predicted elements yet to be discovered. Several of his predictions, for scandium (Sc, 21), gallium (Ga, 31), germanium (Ge, 32), technetium (Tc, 43), and protactinium (Pa, 91), accurately described the chemical properties of these undiscovered elements.

Grouping Methods
The Periodic Table is arranged by rows, also known as periods, and columns, which are also known as groups. Periods are arranged by atomic number while groups are ordered by similar chemical properties. For example, the alkali metal group, which contain sodium and potassium, are all shiny, soft, highly reactive, and have one valence electron in their outermost s-orbital.

Groups
Groups are ordered from 1 to 18 (New IUPAC numbering). Elements in the same group usually have similar chemical properties.

Some groups have names, which are dependent on chemical properties.


 * The alkali metals are extremely light, reactive, and soft metals. Some spontaneously burn in air, and they all react violently on contact with water.
 * Alkaline earth metals are slightly less reactive than the alkali metals, but are still reactive and flammable, but less so than their neighbors.
 * Transition metals are what most people think of when they hear the word "metal." They are generally tough, stable, and hard to melt, with some exceptions. Technetium (Tc, 43) is radioactive, along with all the group 7 transition metals, and mercury (Hg, 80) is liquid at room temperature.
 * The coinage metals are a subgroup in the transition metals. They are very good conductors of electricity, and silver (Ag, 47) is the best conductor of any element.
 * The volatile metals are also a subgroup of the transition metals. They have lower melting points than the other transition metals, and its fourth member, the highly radioactive element opernicium, is even predicted to be a metal gas.
 * The boron group is generally reactive, but properties vary by element, although all can react as +3. As you go down the column, the likelihood of reacting as +1 increases.
 * Properties of the carbon group elements vary widely.
 * The pnictogens'  properties vary widely as well.
 * Like the boron group, the chalcogens are generally moderately reactive, but have variable properties.
 * The halogens are all very reactive, because they have seven electrons in their outer shell and desperately need eight. Fluorine (F, 9) is the most reactive of all and can even form compounds with noble gases.
 * The noble gases are the most inert group of all. Under normal circumstances, they will not form compounds. However, in 1962, scientists made xenon (Xe, 54) compounds with ultra-reactive fluorine (F, 9).

Special Blocks
There are also several blocks of elements that have special names.


 * The lanthanides and the actinides are the elements at the very bottom of the condensed, 18-column form of the Periodic Table.


 * The lanthanides, sometimes known as lanthanoids, are elements 57-71


 * The actinides, sometimes known as actinoids, are elements 89-103


 * The lanthanides and actinides, along with scandium (Sc, 21) and yttrium (Y, 39) are collectively known as rare earth elements or inner transition metals based on the incorrect assumption that they were extremely rare due to the ineffective separation techniques used to isolate the rare earths, which are commonly found in the same ores. While this is true for several of the rare earths, such as protactinium (Pa, 91), most rare earths are actually quite common.


 * The lanthanides' general purposes involve their dynamic magnetic and optical properties, as they are components of strong magnets, such as the famous neodymium-iron-boron type. The lanthanides are also commonly used in lasers. However, cerium has uses for the fact that it can create sparks, europium is used in computers screens, and terbium's ability to change shape in a magnetic field allows the element to turn a table into a loudspeaker.


 * As for the actinides, most of their purposes are not beyond experimental. However, thorium is used for purposes similar to transition metals (its half-life in 14 000 000 000 years), uranium-235 is used in nuclear reactors and weapons, uranium-238 is used in old merchandise (from the radiation-craze era), plutonium is used in weapons and old pacemakers, americium in smoke detectors, and californium in some neutron generators.
 * Elements past uranium are known as transuranic, and do not occur in large quantities in nature, although neptunium and plutonium do occur in traces in uranium and thorium ore due to side reactions triggered by the spontaneous fission of uranium and thorium, and several other actinides have been detected in the spectra of stars.


 * Boron, Silicon, Germanium, Arsenic, Antimony, and Tellurium are known as semi-conductors, because they can conduct electricity under certain conditions - that's why electronics use silicon.


 * With Polonium, these seven elements are known as metalloids - they have some metallic properties and some nonmetallic properties. However, some are more metallic than others, and these are to the left of this "divide" that starts just below boron and "stair-steps" down the periodic table towards the lower righthand corner. Elements to the left are more metallic; elements to the right are more nonmetallic.


 * Elements in Groups 13-16 that are not categorized are just simply known as "other metals/non-metals." Hydrogen is also considered in the"other nonmetals."

Some examples of Periodic Trends

 * Atomic Radius: distance from nucleus to outermost electron; increases as one moves down a period, and from right to left. These are due to more electron shells and lower electronegativity - except in the case of the noble gases, which have enough electrons already and so have no electronegativity, but the protons have enough electromagnetivity to pull electrons in.
 * Ionization Energy: amount of energy required to move one electron from an atom; increases as one moves up a period, and from left to right
 * Redox properties: the possibility the element will be involved in a redox equation; increases as one moves outward from the center of the table, with the exception of the inert noble gases.

Other periodic trends are only applicable for specific groups. For example, the melting points of the alkali metals decrease and their reactivity increases going down the group while the melting points of the halogens increase and their reactivity decreases when similarly going down the group.

Bonding Trends
The location of an element on the Periodic Table governs the element’s bonding behaviors. For example, sodium (Na, 23) is an alkali metal, so it prefers to give away its outermost electron and form salts, while neon (Ne, 10), being a noble gas, refuses to bond due to its full outer shell.

Ionic vs. Covalent
An ionic bond is a chemical bond where one ion loses one or more electrons and transfers them to another ion. These are generally seen in metal-nonmetal combinations, such as sodium chloride (NaCl), in which an electron is transferred from the sodium atom to the chlorine atom.

A covalent bond is a chemical bond where one or more electrons from one element leave that element and fill the outermost electron shell of the other element. In most simple cases, this results in both elements having full outer shells of electrons. These are usually seen in nonmetal-nonmetal bonds, such as in water (H2O), in which the electrons are shared between the hydrogen atoms and the oxygen atom.

Ionic Charges
If an atom gains or loses electrons, it is known as an ion. When it loses electrons, it has a positive charge and is known as a cadion. When an ion gains electrons, it has a negative charge and is called an anion.

Metallic bonds
When two or more metals form a compound, the resulting bond is known as a metallic bond. It is the metallic equivalent of a covalent bond.

Polar covalent bonds
When the atoms in a covalent bond do not share the electrons equally, it is called a polar covalent bond. For example, in water (H2O), the oxygen atom tends to tug on the shared electrons more than the hydrogen atoms, thus creating a slightly negative charge around the oxygen atom.

Solubility Trends
For information about solubility trends, please see Chem Lab/Aqueous Solutions.

Links

 * Chemistry and Mineral Terms section on GMOA Notes
 * Nice explanation of a few periodic properties
 * [[Media:Inf flat Periodicity notes.docx|Infinity Flat's Periodicity Notes]]