The Nomenclature of Inorganic Compounds: The Periodic Table of Elements

Discussion no. 2 The Nomenclature of Inorganic Compounds
Discussion no. 3
The Periodic Table of Elements

Sta. Maria, Yza
Suaco, Trisha Marie T.
Teves, Joan Marie
Vale, Kate
IN-9 Group 9

I. Introduction
Representing molecules as images allows us to impart a great deal of information concerning molecular structure. But these molecules can also be named, and there are occasions when this is more convenient than drawing a picture. In the early days of chemistry, the list of known compounds was short that chemists could memorise the names of all of them. Some of the unique names given by chemists in the early days were salt, cinnabar and laughing gas – names which tells us nothing about what the compound is.
New compounds were often named by their origin of place, their discoverer, their physical appearance or properties. Today, more than 20 million compounds are known and thousands more are synthesized each year. Chemists consequently need systematic procedures for naming chemical compounds. The International Union of Pure and Applied Chemistry (IUPAC) has established uniform guidelines for naming various types of chemical substances. II. Discussion

The Nomenclature of Inorganic Compounds
Objectives: To be familiar with the different rules governing the formulation of a compound’s chemical name as well as its formula. Chemical compounds are divided into two groups and these are: Inorganic and Organic. Organic compounds are composed of principally carbon, hydrogen and oxygen (these are the compounds of life) while all other compounds are called inorganic compounds.
Classification of Compounds: A. Binary Compounds 1. Metal and Non-metal
Inorganic compounds composed of just two elements, metal cations (positive ions) and non-metal anions (negative ions). These are referred to as metal-nonmetal binary compounds. There are two types of metals, some metals form cations with only one charge (e.g.Na+, K+) while other cations form two or more charges (e.g.Ca2+, Al3+).
Metals with Fixed Oxidation Number
Group IIA ( Alkaline Earth Metals) (Zn &Cd- transition metals)
Group IIA ( Alkaline Earth Metals) (Zn &Cd- transition metals)
Group IA (Alkali Metals) (Ag- transition metals)
Group IA (Alkali Metals) (Ag- transition metals)

Cations | Name | Li+ | Lithium | Na+ | Sodium | K+ | Potassium | Rb+ | Rubidium | Cs+ | Cesium | Ag+ | Silver | Cations | Name | Be2+ | Beryllium | Mg2+ | Magnesium | Ca2+ | Calcium | Ba2+ | Barium | Zn2+ | Zinc | Cd2+ | Cadmiun |

Cations | Name | Al3+ | Aluminum |
Group IIIA
Group IIIA

GUIDELINES FOR BINARY IONIC COMPOUNDS 1. The cation is listed first before the anion 2. The cation-anion must give a net charge of zero 3. The ratio is always expressed by the smallest whole numbers 4. The unchanged name of the metal is used. 5. The name of the anion includes only the English root plus –ide
Criss-Cross Method- The numerical value of the charge on the cation becomes the subscript on the anion and vice versa. Example | Criss-Cross Method | Sodium Sulfide | Na+ S2- = Na2S | 2(+1)+(-2)= 0 | Aluminum Chloride | Al3+ Cl- = AlCl3 | (+3)+3(-1)= 0 | Example: Metals with Varied Oxidation Number

* 1.2.1 Stock Method

Stock Method - the name of the metal is followed immediately by a Roman Numeral in parentheses to indicate the charge of the ion. (e.g. Iron (III), Arsenic (IV)) .

Using The Stock Method
In using the Stock Method, one will use a simple algebra equation to determine the charge on the metal and the proper roman numeral to use for more complex compounds.
[(Number of Metal Cations)X(charge on metal)] + [(number of non-metal cations)X(-charge on nonmetal)] = 0
Example:
1.) Name of the compound Fe2S3
[(2x(Fe Charge)] + [3x (S charge)] = 0
2 Fe + (3 x -2)= 0
2 Fe = +6
Fe= +3 (III)
The compound is therefore Iron (III) Sulfide
2.) Finding the formula for Lead (IV) Oxide
Pb4+ O2
= Pb2O4 = PbO2
Explanation : The answer was PbO2 because the Ratio must be expressed by the smallest whole number ratio.

Metals with Variable Oxidation Number using the Stock Method Ion | Stock Name | As3+ | Arsenic (III) | As5+ | Arsenic (IV) | Au+ | Gold (I) | Au3+ | Gold (III) | Cos2+ | Cobalt (II) | Co3+ | Cobalt (III) | Cu+ | Copper (I) | Cu2+ | Copper (II) | Cr2+ | Chromium (II) | Cr3+ | Chromnium (III) | Fe2+ | Iron (II) | Fe3+ | Iron (III) | | Ion | Stock Name | Hg22+ | Mercury (I) | Hg2+ | Mercury (II) | Mn2+ | Manganese (II) | Mn3+ | Manganese (III) | Ni2+ | Nickel (II) | Ni3+ | Nickel (III) | Pb2+ | Lead (II) | Pb4+ | Lead(IV) | Pt2+ | Platinum (II) | Pt4+ | Platinum (IV) | Sb3+ | Antimony (III) | Sb5+ | Antimony (V) | Sn2+ | Tin (II) | Pt4+ | Tin (IV) | |

* 1.2.2 Classical Method

Classical Method- identifies the ion with lower charge number by the Latin name of the element followed with an –ous ending and the higher charged number ends in –ic. If the symbol of an element is derived from the Latin word, the Latin root is generally used rather than the English root.(e.g Arsenic, Chromous, Cupric)

Metals with Variable Oxidation Number using the Stock Method

Ion | Classical Name | As3+ | Arsenous | As5+ | Arsenic | Au+ | Aurous | Au3+ | Auric | Co2+ | Cobaltous | Co3+ | Cobaltic | Cu+ | Cuprous | Cu2+ | Cupric | Cr2+ | Chromous | Cr3+ | Chromic | Fe2+ | Ferrous | Fe3+ | Ferric | | Ion | Classical Name | Hg22+ | Mercurous | Hg2+ | Mercuric | Mn2+ | Manganous | Mn3+ | Manganic | Ni2+ | Nickelous | Ni3+ | Nickelic | Pb2+ | Plumbous | Pb4+ | Plumbic | Pt2+ | Platinous | Pt4+ | Platinic | Sb3+ | Antimonous | Sb5+ | Antimonic | Sn2+ | Stannous | Pt4+ | Stannic | |

2. Two Non-metals A molecular compound is formed when both compounds are non-metals. The least electronegative atom is written first then next by the more electronegative elements wherein the suffix “–ide” is added to the name of the higher electronegative ion. The Greek prefixes are used to indicate the number of atoms of each element that are present in a compound. Prefix | Meaning | mono- | 1 | di- | 2 | tri- | 3 | tetra- | 4 | penta- | 5 | hexa- | 6 | hepta- | 7 | octa- | 8 | nona- | 9 | deca- | 10 |

The prefix –mono is rarely used because when no prefix appears, it is assumed that only one atom is used and when the two vowels appear next to each other like “oo” in monoxide, in which the vowel from the prefix is dropped.

Examples: a. N2O4 = dinitrogen tetroxide b. Cl2O7 = dicholrine heptoxide c. N4S4 = tetranitrogen tetrasulfide d. P4S10 = tetraphosphorus decasulfide 3. Hydrogen and A Non-metal

3.1 Aqueous State & 3.2 Gaseous State Acid is a substance that produces hydrogen ions in water. Binary Acids are consist of two elements one is a hydrogen-contaning compound. Pure compounds (ie- in the gaseous state), have different properties than the acqueous state of the compound. The rule in naming a pure compound is different to naming the acid in aqueous solution.
Rules in Naming Binary Acids
Binary Acids in Gaseous State (ie- in the gaseous or liquid state) 1. Start with “Hydrogen” 2. Followed by the anion name ending ‘-ide’
e.g. Hydrogen Chloride
Binary Acids in Aqueous State 1. Start with the prefix “hydro-”. 2. Change the anion name at the end with the suffix “-ic”. 3. Then add the word “acid” at the end.
e.g. Hydrochloric Acid Anions | Formula (Binary) | Gasseous State | Aqueous State | F- | HF | Hydrogen Flouride | Hydroflouric Acid | Cl- | HCl | Hydrogen Chloride | Hydrochloric Acid | Br- | HBr | Hydrogen Bromide | Hydrobromic Acid | I- | HI | Hydrogen Iodide | Hydroiodic Acid | S- | H2S | Hydrogen Sulfide | Hydrosulfuric Acid | Hydroxide and a Metal
Hydroxides are ternary compounds containing a metallic element. They are ionic compounds formed by metal cations and anions OH-.
How are they named?
It is formed by two words. The first word is the name of the metal, followed immediately by the oxidation number between brackets and in Roman numbers. If the oxidation number of the metal is invariable, it is not indicated. The second word is ‘hydroxide’. The oxidation number of the metal is equal to the number of ions OH-. For example: Cr(OH)3 this would be named as Chromium(III) hydroxide
Remember that you must know the symbols of the elements and the oxidation numbers which are invariable. Formula | Name | LiOHBa(OH)2Fe(OH)3Cr(OH)3Al(OH)3 | Lithium hydroxideBarium hydroxideIron(III) hydroxideChromium(III) hydroxideAluminum hydroxide |

The Hydroxides are similar to oxides. In their case the metals are combined with the hydroxyl ion (OH)- rather than O2-. Hydroxides easily dissociate on heating. The (OH)- ions are driven off as water leaving behind the metal oxide. The structure of Hydroxides is strongly dependent on the radii of the constituent ions.
The Oxides and Hydroxides are divided into the following groups: * Periclase Group - includes all oxides that have bivalent (two positive charges) metallic ions and the halite structure. In this arrangement each metallic ion is shielded by siz oxygen ions around it. Only metallic ions of medium size adopt this structure. Smaller ions are surrounded by four oxygen ions and larger ones by six. In oxides where the bonding is more covalent the crystal structures are determined by both the configuration of the electrons and the size of the ion. Periclase minerals, like most oxides, are chemically reactive and therefore found in unusual environments. For example Manganosite MnO forms in marbles from the breakdown of dolomite. Amongst minerals in this group only Periclase is common. * Zincite Group - includes Zincite and Bromellite. Both are rare. The Zn2+ or Be2+ are small enough to be adequately shielded by four oxygen ions. Each oxygen in turn is surrounded by four Zn or Be ions. The resulting hexagonal crystal adopts a hemimorphic habit. In other words, the two terrminations have differing faces, indicating a lacking centre of symmetry. Bromellite is much harder than Zincite because the smaller Be ion forms a more strongly covalent bond with its neighboring oxygens. * Cuprite Group - in this group the metal is surrounded by only two oxygen ions. Minerals, including Cuprite, crystallise in the cubic system, often also forming octahedra, dodecahedra or a combination of all three. Because the bonds with oxygen have a large degree of covalency the minerals tend to be brittle. * Corundum Group - also commonly called the Hematite Group. Includes the simple oxides Corundum (Al2O3), Hematitie (Fe2O3) and other oxides with the general formula ABO3. All have hexagonal structures with the metal surrounded by six oxygen ions. The metals may be trivalent (eg. Al3+) or may be a mixture of bivalent and tetravalent metals such as Fe2+ and Ti4+, as in Ilmenite. The possibility of metal substitution means that the group is prone to form solid solutions . An example is the Ilmenite- Pyrophanite series. Members of the group typically form at high temperatures in igneous rocks or during high-temperature metamorphism of silica-deficient rocks. Hematite and Corundum are the only common species. * Spinel Group - probably the largest and most complex group of oxides. Many exist as members of several solid solution series. They are mixed oxides containing a combination of metals with a general formula AB2O4. Some metals are stabilised by four oxygen atoms in a tetraherdal structure whilst others are stabilised by six in an octahedral arrangement. Still others can occur in both these positions. The oxygen ions are held in a cubic close packed arrangment with the metals. The combinations gives rise to very complex crystal lattice structures. Most minerals in the group form at high temperature in igneous and metamorphic rocks, usually those lacking silica. * Rutile Group - includes all oxides with the general formula MO2, in which the metal is tetravalent (carries four positive charges). The metal is surrounded by six oxygen ions, giving rise to the rutile structure. Typically the metals involved are titanium, manganese, tin and lead. However, because of the significant differences between these metals their oxides do not form solid solutions and the minerals have widely different origins. Brookite, Anatase and Rutile are polymorphs of titanium dioxide (TiO2) having the same composition but differening arrangement of ions in the rutile structure. * Uraninite Group - members include Uraninite and Thorianite, oxides of uranium and thorium respectively. They have cubic structures resembling Fluorite with each metal ion surrounded by eight oxygen ions. Because of the chemical and physical similarity of the metals they form a complete solid solution series, producing a wide range of minerals of intermediate composition. Nearly all crystallise in the cubic system as octahedra, cubes or combinations. * Diaspore Group - includes oxyhydroxides of trivalent metals including Al3+, Fe3+ and Mn3+. The general chemical formula is MO(OH). Each metal ion is surrounded by six negative ions, three O2- and three (OH)-. Group minerals are important sources of key industrial metals and are associated with other hydrous minerals. Some, like Bauxite, do not form crystals but occur as noncrystalline colloidal precipitates. Bauxite is usually the result of prolonged weathering of aluminous rocks where the silica content has dissolved leaving behind the aluminium hydroxide as residue. * Brucite Group - includes the hydroxides of divalent metals including Magnesium Mg2+ and Manganese Mn2+. The arrangement is octahedral with the metal at the centre. The structure is usually layered, consisting of six hydroxyl (OH)- ions surrounding the metal ion. The layers are stacked upon each other and held together by weak hydrogen bonds. The minerals therefore adopt plate-like habits, have very pronounced cleavage and are very soft.
TERNARY COMPOUNDS
An ionic compound is named using the name of the cation followed by the name of the anion, eliminating the word ion from each. Unlike the naming of molecular compounds, no Greek prefixes are used.
Ternary compounds contain atoms of three different elements. These compounds usually consist of metal ion or hydrogen and a negatively charged polyatomic ion. It is necessary to know the names of the polyatomic ions in order to name ternary compounds.
In naming and writing ternary compounds, the same rules as for binary compounds are followed. The only difference is that the name of the polyatomic ion replaces the name of the anion. As with ionic compounds, no prefixes are used. Consider the compounds below: K2CO3 potassium carbonate Fe2(SO4)3 iron (III) sulfate Zn3(PO4)2 zinc phosphate
Notice that parentheses are used around a polyatomic ion when more than one anion is needed in the formula as in iron (III) sulphate and zinc phosphate.

Compounds with Polyatomic Ions | Formula | Name | Use | (NH4)2SO4CaCO3FeSO4PbCrO4Li2CO3Mg(OH)2NaHCO3NaNO2KNO3ZnSO4 | ammonium sulfatecalcium carbonateiron(II) sulphatelead(II) chromatelithium carbonatemagnesium hydroxidesodium bicarbonatesodium nitritepotassium nitratezinc sulfate | fertilizercalcium supplement, chalkiron supplementartists’ yellow oil painttreat manic depressionantacidantacid, baking sodameat preservativeplant bloom boostertissue repair, healing of wounds |

Metals with Polyatomic Ions
Ions that consist of a combination of two or more atoms are called polyatomic ions. Polyatomic ions are composed of more than one atom such as the cation formed from non-metal ions and hydrogen atoms, with names that end in –ium.
For example H3O- hydronium ion NH4+ ammonium ion
Most polyatomic anions containing oxygen (oxyanions) end in –ate or –ite. Common Polyatomic Ions and Their Charges | Formula | Name | AsO43-AsO33-BO33-
AlO33-P2O74-As2O74-SiO44- | arsenatearseniteboratealuminatepyrophosphatepyroarsenateorthosilicate |

Oxyacids and Oxyanions
Oxyanions or oxoanions are polyatomic anions that contain one or more oxygen atom and one atom (the ‘central atom’) of another element. Often, two or more oxyanions have the same central atom but different numbers of O atoms. Starting with the oxyanions whose names end in –ate, we can name these ions as follows: 1. The ion with one or more O atom than the –ate ion is called the per … ate ion. 2. The ion with one less O atom than the –ate ion is called the –ite ion. 3. The ion with two fewer atoms O atoms than the –ate ion is called the hypo … ite ion.
There is another important class of acids known as oxyacids or oxoacids, which ionize to produce hydrogen ions and the corresponding oxyanions. The formula of an oxyacid can be determined by adding enough H+ ions to the corresponding oxyanion to yield a formula with no net charge. For example, the formulas of oxyacids based on the nitrate (NO3-) and sulphate (SO42-) ions are HNO3 and H2SO4, respectively. The names of oxyacids are derived from the names of the corresponding oxyanions using the following guidelines: 1. An acid based on an –ate ion is called … ic acid. 2. An acid based on an –ite ion is called … ous acid. 3. Prefixes in oxoanion names are retained in the names of the corresponding oxoacids.
Many oxyacids, such as H2SO4 and H3PO4, are polyprotic—meaning that they have more than one ionizable hydrogen atom. In these cases, the names of anions in which one or more (but not all) of the hydrogen ions have been removed must indicate the number of H ions that remain, as shown for the anions deprived from phosphoric acid:
H3PO4 phosphoric acid
H2PO4- dihydrogen phosphate ion
HPO42- hydrogen phosphate ion
PO43- phosphate ion Formula Name | Formula Name | CationsNH4+ ammoniumH3O+ hydroniumHg22+ mercury (I)Diatomic anionsOH- hydroxideCN- cyanideAnions with carbonCO32- carbonateHCO3- hydrogen carbonateCH3COO- acetateC2O42- oxalate | OxyanionsSO42- sulfateSO32- sulfiteNO3- nitrateNO2- nitritePO43- phosphateMnO4- permanganateCrO42- chromateCr2O72- dichromateClO- hypochlorateClO2- chloriteClO3- chlorate |

III. References * Exploring life through Science : Chemistry (2007) by Aristea V. Bayquen * Introductory Chemistry (1998) by Darrel D. Ebbing * Introductory Chemistry for Today (2008) by Spencer C. Seager * Fundamentals of Chemistry (1999 ) by Ralph A. Burn * Breaking Through Chemistry (2006) by Baguio Buturan * Chemistry 2nd Edition (2009) by Burdge * CHEMISTRY (2001) by Blackman, Bottle, Schmid, Mocerino and Wille * General Chemistry 2nd Edition (2003) by John B. Russel

THE PERIODIC TABLE OF ELEMENTS

Introduction: The periodic table tells us all forms of chemical and physical information about metals, nonmetals, and metalloids as well as their compounds. It shows us the systematic arrangement of the elements and later on compares their similarities and differences. The need for organization was recognized by early chemists and there were numerous attempts to discover relationships among the properties of elements. As the number of known elements increased, chemists developed the periodic table to organize and classify these elements. The development of the periodic table brought system and order. It also helped the chemist to predict the elements that had to be discovered in the future.
Objectives:
All the end of the laboratory discussion, the students should be able to understand the arrangement of the periodic table into periods and groups; known the names, symbols, and some physical properties. Correlate an element’s position in the periodic table with its atomic structure; and compare the periodic trends existing between elements in the periodic table.

A. Brief History of the Periodic Table
The periodic table was formed by scientists before who were trying to classify the elements according to their properties. In 1829, German chemist, Johann W. Dobereiner discovered triad. This is made up of 3 elements where in the middle element had an atomic mass almost equal to the averages atomic mass of the other two elements. However, it cannot be explained. In 1864, English chemist, John Newlands observed that the elements are arranged in the order of increasing atomic masses, there appeared to be a repetition of similar properties for every 8th element. He called it low of octaves. However, after the element calcium, this is not applicable anymore. In 1864, German chemist, Julius Lothar Meyer devised a table which has periodic variations in properties and has 56 elements. Also in 1864, Russian chemist, Dmitri Ivanovich Mendeleev arranged the elements according to increasing atomic weight but there are still some irregularities. (Ex. “Te and I”, if we will follow the increasing atomic weights, “I” must be in the left of ”Te”. But to follow similar characteristics, “I” must be in line with “Cl, Br and F”. then it will be in the right of ”Te”) In 1913, English physicist, Henry H. Moseley used atomic numbers to arrange the periodic table. And until now, this table is what we use in the modern times.

Parts of the Periodic Table * Elements
Elements are pure substances that are made up of only one type of atom. * Atomic Number and Atomic Mass
As the table progresses from left to right, and top to bottom, the atomic number of the elements increases. The atomic number is the count of the protons contained in the atomic nucleus. The table also shows atomic mass, the total mass of the neutrons, protons and electrons that make up the atom. For elements with no stable isotope, the table gives in parenthesis the atomic mass of the isotope with the longest half-life; in other words, the most stable form of the element.

* Group or Family
The table 's 18 columns, read vertically from top to bottom, represent groups. All elements in a group have the same number of electrons orbiting the nucleus in the outermost shell. The exceptions to this rule include hydrogen, helium and the "transitional elements," which occupy groups three through 12. Elements in a group share important chemical characteristics. Group 18, for example, includes the "inert" or "noble" gases. Group 17 includes the five halogens. * Period
The seven rows in the table represent periods. Each element in a single row has the same number of electron shells, which surround the atomic nucleus. The elements hydrogen and helium have a single orbital shell; elements in the second row have two orbitals, and so on. In the seventh period, elements have a seventh orbital shell, which is the highest energy level occupied by naturally occurring electrons. * Graphic Indicators
Some periodic tables display a color code that shows the state of the element--solid, liquid, gas, or unknown--at zero degrees Centigrade. Borders may show whether the element is naturally occurring (solid border), occurring only as a result of radioactive decay (dashed border) or artificial (dotted border). A single thick line sometimes appears in the periodic table dividing the elements into metallic (to the left) and non-metallic (to the right). * Lanthanides and Actinides
At the bottom of the periodic table are two additional rows of 14 elements each. The top row shows the lanthanides, elements 58 through 71. The bottom row is the actinides, elements 90 through 103. The first elements in these two series are contained in the main body of the periodic table: lanthanum (57) and actinium (89). These are the 30 rare earth metals, most of which are synthetic.

The Element Groups * Group IA: The Alkali Metals
The alkali metals lithium (L), sodium (Na), potassium (K), rubidium (Rb), Cesium (Cs), and francium (Fr) are silver gray, soft metals that can be cut with the knife. They have very low densities and are good conductors of heat and electricity. Group IA metals are quick to react with water, oxygen, and other chemicals and are never found as free (uncombined) elements in nature. Typical alkali metals compounds are water soluble and are present in seawater and salt deposit. Because these metals quickly react with oxygen, they are sold in evacuated containers but, once opened; they usually are stored under mineral oil or kerosene. * Group IIA: The Alkaline Earth Metals
Group IIA metals include beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). These metals have melting points that are higher than the Group IA metals. Their densities are all low but are still somewhat higher than densities of comparable alkali metals. They are less reactive than the alkali metals. All alkaline earth metals have two valence electrons and form ions with a positive two (2+) charge. * Group IIIA
The first element in Group IIIA is boron (B), a metalloid with a very high melting point and predominantly nonmetallic properties. The other in this group include aluminum (Al), gallium (Ga), indium (In), and thallium (TI), which form ions with a positive three (3+) charge. Density and metallic character increases as atomic number increases within this group. * Group IVA: The Carbon Family
Metallic character increases from top to bottom for the Group IVA elements carbon (C), silicon (Si), germanium (Ge), thin (Sn) and lead (Pb). Differences in the crystal line arrangements of carbon atoms account for the hardness of diamond and the slipperiness of black graphite.

* Group VA
Elements in Group VA include the nonmetals nitrogen (N) and phosphorus (P), the metalloids arsenic (As) and antimony (Sb), and the heavy metal bismuth (Bi). Thus, there is a dramatic change in appearance and properties from top to bottom in this group. * Group VIA
These group elements, called the oxygen family, include oxygen (O), sulfur (S), selenium (Se), tellurium (Te), and polonium (Po). Although they all have six valence electrons, properties vary from nonmetallic to somewhat metallic as atomic number increases. * Group VIIA: The halogens
Group VIIA, the halogen family, includes fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At). The family name, halogen, comes from Greek words meaning “salt former”. Each halogen atom has seven valence electrons. As elements, the halogens are all diatomic; they have two atoms per molecule. Halogens are too reactive to be found free in nature. * Group VIIIA: The Noble Gases
Elements in Group VIIIA at the far right of the periodic table are known as the noble gases. This family includes helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). Noble gases all exist as monatomic (single) gaseous atoms that do not tend to enter into reactions with other elements.

Classification of Elements in the Periodic Table * Metals
Is a material that has a characteristic luster or shine and that is a relatively good conductor of heat and electricity. Metals have high densities and melting temperatures. They are lustrous and are excellent conductors of heat and electricity. Metals are usually solid and their oxides are often soluble in acidic solutions. They are ductile which means they can be shaped into wires. Metals are also malleable, meaning they can be rolled up or hammered into sheets. * Non-metals
Nonmetals tend to be brittle when they exist as solids. They exhibit no metallic luster. They have low densities and are poor conductors of heat and electricity. Most non metals have low melting points, and many exist as gases at room temperature. * Metalloids
Semiconductors or metalloids, such as the elements silicon (Si), germanium (Ge), and arsenic (As), exhibit both metallic and non metallic properties. Like nonmetals, they are able to conduct electricity under certain conditions.

* Transition Metals
The transition metals are located in the central region of the periodic table and are readily identified by a Roman numeral followed by the letter “B” on many tables. In general, the properties of the transition metals are rather similar. These metals are more brittle, they are harder, and they have higher melting points and boiling points than do other metals. * Inner Transition Metals
Inner transition metals are the two rows of elements at the bottom of the periodic table. Locate lanthanum with atomic number 57 in the periodic table.

Electronic Configuration using the Periodic Table
The electron configuration for an element is the arrangement of electrons in the orbits (or shells) of a neutral atom. Shells closer to the nucleus have higher binding energy.

* Valence shell electrons
The outmost shell occupied by electrons in an atom. Valence shell electrons are the outer electrons of an atom that are involved in chemical bonding. * Sublevel
The sublevel occupied by any electron is determined by the electron 's angular momentum quantum number, l. It is found by solving the Schrodinger equation, allowing us to find probability distributions for an electron in an atom. For example, we can say that an electron in a hydrogen atom 's 1s sublevel will be found 99 percent of the time somewhere in a sphere with a given radius around the nucleus. Hence the reason we can draw the s sublevel as a sphere. Electron orbital designated s, p, d or f. These sublevels or orbitals have characteristic shapes which can be used to explain and predict the chemical bonds that atoms can form. s sublevels are spherically shaped. The p, d and f sublevels have more complex shapes.

* S-block elements
The s-block of the periodic table consists of the first two groups, namely the alkali metals and the alkaline earth metals. The elements in the s-block generally exhibit well-defined trends in their physical and chemical properties, changing steadily moving down the groups. Their properties can be most readily explained in terms of their electron configuration, with their valence electrons occupying s-orbitals. By this definition, hydrogen and helium are sometimes also considered to be part of the s-block. * P-block Elements
The p-block of the periodic table of the elements consists of the last six groups except helium (which is located in the s-block). In the elemental form of the p-block elements, the highest energy electron occupies a p-orbital. The p-block contains all of the nonmetals (except for hydrogen and helium which are in the s-block) and semimetals, as well as the post-transition metals. * D-block Elements
Any element with an atom that has an incomplete d sublevels or that forms a cation or cations with incomplete d sublevels. * F-block Elements
The elements in both the lanthanide and actinide series (inner transition elements) have f orbitals that are incompletely filled.

*The notation seen on period tables for electron configurations uses the form: nOe n is the energy level
O is the orbital type (s, p, d, or f) e is the number of electrons in that orbital shell. * Aufbau Principle

Is used to determine the electron configuration of an atom, molecule or ion. The principle postulates a hypothetical process in which an atom is "built up" by progressively adding electrons. As they are added, they assume their most stable conditions (electron orbitals) with respect to the nucleus and those electrons already there.

* Hund’s Rule are used to determine the term symbol that corresponds to the ground state of a multi-electron atom. In chemistry, rule #1 is especially important and is often referred to as simply Hund 's Rule.

PERIODIC TRENDS Periodic trends are specific patterns that are present in the periodic table, which illustrate different aspects of a certain element, including its size and its properties with electrons. The main periodic trends include: electronegativity, ionization energy, electron affinity, atomic size (atomic radius), ionic size, metallic property and non-metallic property.

A. ATOMIC SIZE (ATOMIC RADIUS)
The size of an individual atom is difficult to measure since the electron cloud surrounding the nucleus has no sharp boundary. However, it is possible to measure the radius of an atom. The radius of an atom is determined by measuring the distance between the nuclei of two atoms. The distance between the nuclei of two atoms bounded together is called interatomic bonded distance. The atomic radius is defined as half the distance between the nuclei of two like atoms joined together by single bond. * Atomic size generally increases as you move down a group of the periodic table. This is because as you move down a group, the number of energy levels or shells increases. * The size of atoms in a period generally decreases from left to right across the row. This is because atomic radius decreases due to increasing nuclear charge. As you move from left to right across a period, the atom’s nuclei gain more protons. Although there is a gain of protons, there is also an accompanying gain of electrons.
Within each group (or family) of elements, the size of atoms increases each time a new higher energy level is occupied by more electrons. However, within each period of elements, the size of atoms tends to decrease as more electrons are added to a specific energy level. This is because each element in a period has one more proton than the previous element, and the increase I positive nuclear charge draws the electron cloud closer to the nucleus.

B. IONIC SIZE
The ionic radius is the measure of an atom 's ion in a crystal lattice. Values for ionic radius are difficult to obtain and tend to depend on the method used to measure the size of the ion. A typical value for an ionic radius would be from 30 pm to 200 pm. Ionic radius may be measured using x-ray crystallography or similar techniques. * As you move across a row of period of the periodic table, the ionic radius decreases for metals forming cations, as the metals lose their outer electron orbitals. The ionic radius increases for nonmetals as the effective nuclear charge decreases due to the number of electrons exceeding the number of protons.

C. ELECTRON AFFINITY Electron affinity is a measure of the energy charge that occurs when a gaseous atom gains an electron. The electron affinity of an element is the energy change when an electron is added to a gaseous atom to form a 1- ion. As the term implies, the electron affinity is a measure of the attraction an atom has for an additional electron. The amount of energy released when a gaseous atom gains an electron is called the electron affinity. The more positive an atom’s electron affinity, the greater its tendency to accept an electron or electrons, thereby forming anions. * Electron affinity increases across a period from left to right. The halogens possess the greatest electron affinities. When a halogen accepts an electron, it obtains the stable electron configuration of the noble gas to its right. * As one goes down a group, electron affinity generally decreases. This affinity decreases occurs because the valence electrons in the higher period elements are farther from the nucleus and therefor, not much energy is released when an electron is accepted into the valence shell.

D. ELECTRONEGATIVITY Electronegativity is the ability of an atom in a compound to attract additional electrons towards itself. The greater the electronegativity, the greater the attraction for electrons. Small atoms like the elements from Groups VIIA which have outer shell that are nearly filled, attract electrons more easily that larger ones and therefore, tend to have higher electronegativity values. * Electronegativity tends to increase from left to right across each period and deceases from top to bottom in each group. The elements in Group IA (alkali metals) have the lowest electronegativity values, while those in Group VIIA (halogens) have the highest values.
The electronegativity generally increases from left to right across a period with the Group VII element having the highest value for the period. The electronegativity generally decreases from top to bottom down a group. Francium is the element with the lowest electronegativity.

E. IONIZATION ENERGY
A certain amount of energy is necessary to knock off the electron from a neutral gaseous atom to form a positive ion called cation. The minimum amount of energy required to do this is called the ionization energy. The greater the ionization energy required, the more tightly held the electron in an atom.
It is an energy required to remove an electron from the outermost shell of an atom, measure of how difficult it is to remove an electron from a gaseous atom. Energy must always be absorbed to bring about ionization, so ionization energies are always positive quantities.
The minimum amount of energy required to dislodge the least firmly attached electron from an atom in the gaseous state is called the first ionization energy. The second ionization energy is the amount of energy required to remove the second electron from the gaseous atom, and so forth. The second ionization energy is always higher than the first, since the second electron being removed is from an inner energy level and it is more tightly held by the nucleus. * The first ionization energy for the atoms going across a period from left to right generally increases. This rise is due to the increasing nuclear charge (atomic number), although the electrons belong to the same principal energy level. Increasing nuclear charge attracts electrons in the same shell more tightly. Ionization energies increase much more gradually across the period for the transition and inner transition elements than for the main group elements. * The ionization energy decreases from one atom to the next as one goes down a group. This energy decrease is due to the fact that the valence electrons at a higher energy levels or at a further distance from the nucleus, which means that the attraction of the electrons to the nucleus becomes smaller. Thus, it is easier to remove those electrons.

F. METALLIC PROPETY; AND NON-METALLIC PROPERTY Elements with three or less than three electrons on the outer energy level are classified as metals. Atoms of metals tend to lose these outer electrons to form positive ions or cations. Elements with five or more electrons on the outer energy level are classified as nonmetal.
The tendency of an element to lose electrons and form positive ions (cations) is called electropositive or metallic character. The tendency of an element to accept electrons to form an anion is called its non-metallic or electronegative character.
In each period, metallic character of elements decreases as we move to the right. Elements to the left have a pronounced metallic character while those to the right have a non-metallic character. Conversely, non-metallic character increases from left to right.
The metallic character decrease from left to right across the period because the elements to the left of the periodic table have a tendency of losing electrons easily as compared to those to the right. As we move from left to right of the period, the electrons of the outer shell experience greater pull of the nucleus. This greater force of attraction is because the nuclear charge increases and the size of the atom decreases from left to right. Thus, electrons of the elements to the right of the table do not lose electrons easily so are non-metallic in nature.
The metallic character increase down the group because as we move down the group the number of shells increases. This causes the effective nuclear charge to decrease due to the outer shells being further away: in effect the atomic size increases. The electrons of the outermost shell experience less nuclear attraction and so can lose electrons easily thus showing increased metallic character. Metallic property generally decreases as you move from left to right across a period and increases from top to bottom within a given group.

References: * Phoenix Science Series Chemistry (2009) by Estrella Mendoza * Introductory Chemistry for Today (2008) by Spencer C. Seager * Chemistry 2nd Edition (2009) by Burdge

References: * Phoenix Science Series Chemistry (2009) by Estrella Mendoza * Introductory Chemistry for Today (2008) by Spencer C. Seager * Chemistry 2nd Edition (2009) by Burdge