The Periodic Table of the Elements is just a way to arrange the elements to show a large amount of information and organization. As you read across the chart from right to left, a line of elements is a Period. As you read down the chart from top to bottom, a line of elements is a Group or Family. We number the elements, beginning with hydrogen, number one, in integers up to the largest number. The integer number in the box with the element symbol is the atomic number of the element and also the number of protons in each atom of the element.
PROPERTIES OF MATTER
The Periodic Table, or Periodic Chart, is based on the properties of matter. A property is a quality or trait or characteristic. We can describe, identify, separate, and classify by properties. How would you describe a person? A young man impressed with a young lady might describe her, “She has long dark hair that she keeps in a pony-tail, brown eyes, a long neck, and a very light complexion. She is about 180 centimeters tall and has pierced ears.” He has used some of her properties to describe her. You might be able to pick her out of a small group of people based on his description if it is not too inaccurate, too vague, or too biased. Similarly, you can collect a number of properties to describe an element or compound. The properties of the element or compound, though, are true for any amount of the material anywhere. South American gold is indistinguishable from South African gold by its properties.
There are two types of property of matter. Physical properties describe the material as it is. Chemical properties describe how a material reacts, with what it reacts, the amount of heat it produces as it reacts, or any other measurable trait that has to do with the combining power of the material. Properties might describe a comparative trait (denser than gold) or a measured trait (17.7 g/cc), a relative trait (17.7 specific gravity), or an entire table of measurements in a table or graph form (the density of the material through a range of temperatures).
Physical properties include such things as: color, brittleness, malleability, ductility, electrical conductivity, density, magnetism, hardness, atomic number, specific heat, heat of vaporization, heat of fusion, crystalline configuration, melting temperature, boiling temperature, heat conductivity, vapor pressure, or tendency to dissolve in various liquids. These are only a few of the possible measurable physical properties.
Chemical properties include: whether a material will react with another material, the rate of reaction with that material, the amount of heat produced by the reaction with the material, at what temperature it will react, in what proportion it reacts, and the valence of elements.
We can separate or purify materials based on the properties. We can separate wheat from chaff by throwing the mix into the wind. The less dense chaff is moved more by the wind than the denser wheat. We can separate a mixture of sand and iron filings by magnetism. The iron filings will stick to a magnet dragged through the mixture. We can separate ethyl alcohol (good old drinking alcohol) from water by boiling point. This process is called distillation. A mixture of water and insoluble material with alcohol mixed in it will release the alcohol as vapor at the boiling point of alcohol (78 °C). We can separate by solubility. A mixture of table salt and sand can be separated by adding water. The salt dissolves and the sand does not.
The periodic chart came about from the idea that we could arrange the elements, originally by atomic weight, in a scheme that would show similarity among groups. The original idea came from noticing how other elements combined with oxygen. Oxygen combines in some way with all the elements except the inert gases. Each atom of oxygen combines with two atoms of any element in Group 1, the elements in the row below lithium. Each atom of oxygen combines one-to-one with any element in Group 2, the elements in the row below beryllium. From here as we investigate the groups from left to right across the Periodic Chart, the story is not quite so clear, but the pattern is there. Group 3 is the group below boron. All of these elements combine with oxygen at the ratio of one-and-a- half to one oxygen. Group 4, beginning with carbon, combines two to one with oxygen. The group of transition elements (numbers 21-30 and 39-48 and 71-80 and 103 up) have never been adequately placed into the original scheme relating to oxygen. The transition elements vary in the ways they can attach to oxygen, but in a manner that is not so readily apparent by the simple scheme. Gallium, element number thirty-one, is the crowning glory of the Periodic Chart as first proposed by Mendeleev. Dmitri Ivanovich Mendeleev first proposed the idea that the elements could be arranged in a periodic fashion. He left a space for gallium below aluminum, naming it eka- aluminum, and predicting the properties of gallium fairly closely. The element was found some years later just as Mendeleev had predicted. Mendeleev also accurately predicted the properties of other elements.
Most Periodic Charts have two rows of fourteen elements below the main body of the chart. These two rows, the Lanthanides and Actinides really should be in the chart from numbers 57 - 70 and from 89 - 102. To show this, there would have to be a gulf of fourteen element spaces between numbers 20 - 21 and numbers 38 - 39. This would make the chart almost twice as long as it is now. The Lanthanides belong to Period 6, and the Actinides belong to Period 7. In basic Chemistry courses you will rarely find much use for any of the Lanthanides or Actinides, with the possible exception of Element #92, Uranium. No element greater than #92 is found in nature. They are all man-made elements, if you would like to call them that. None of the elements greater than #83 have any isotope that is completely stable. This means that all the elements larger than bismuth are naturally radioactive. The Lanthanide elements are so rare that you are not likely to run across them in most beginning chemistry classes. Another oddity of the Periodic Chart is that hydrogen does not really belong to Group I -- or any other group. Despite being over seventy percent of the atoms in the known universe, hydrogen is a unique element.
For more information on each of the elements, see the chapter on the elements alphabetically.
PERIODIC CHART OF THE ELEMENTS
ELEMENT, ION, AND COMPOUND SYMBOLS
For every element there is one and only one upper case letter. There may or may not be a lower case letter with it. When written in chemical equations, we represent the elements by the symbol alone with no charge attached. The seven exceptions to that are the seven elements that are in gaseous form as a diatomic molecule, that is, two atoms of the same element attached to each other. The list of these elements is best memorized. They are: hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine. They are easy to pick out on the periodic table because they (with the exception of hydrogen) form the shape of a seven. The chemical symbols for these diatomic gases are: H2, N2, O2, F2, Cl2, Br2, and I2. Under some conditions oxygen makes a triatomic molecule, ozone, O3. Ozone is not stable, so the oxygen atoms rearrange themselves into the more stable diatomic form.
Chemtutor highly recommends that a few short lists be well learned for immediate recognition. The diatomic gases (hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine), the Group one elements (lithium, sodium, potassium, rubidium, cesium, and francium), the Group two elements (beryllium, magnesium, calcium, strontium, barium, and radium), Group seven elements, the halogens, (fluorine, chlorine, bromine, iodine, and astatine), and the noble gases (helium, neon, argon, krypton, xenon, and radon). If nothing else, learning these as a litany will help you distinguish between radium, a Group 1 element, and radon, an inert gas.
Groups of two or more element symbols attached to each other without any charge on them indicate a compound. CaCl2 is a compound with two chlorine atoms for each calcium atom. CuSO4•5H2O, cupric sulfate pentahydrate, is also a compound. It has one copper atom and one sulfate ion consisting of a sulfur atom and four oxygen atoms attached to five molecules of water.
Charged particles, called ions, when written with symbols will have the charge, either positive (+) or negative (-), written to the right and superscripted to the chemical symbol. For instance, Na+ is the symbol for the sodium ion. Atoms or polyatomic ions with charges of more than one, either positive or negative, have a number with the charge. For instance (CO3)2- is the symbol for the carbonate ion. The carbonate ion has one carbon atom in it, three oxygen atoms, and a charge of negative two. Observe that the charge is outside the parentheses, indicating that the charge is from the polyatomic ion as a whole.
CATEGORIES OF ELEMENTS
What Chemtutor calls 'categories of elements' include; metals, non-metals, semi-metals, noble gases, and hydrogen.
PERIODIC CHART OF THE ELEMENTS
Consider a staircase-shaped line on the Periodic Chart (in blue above) starting between boron and aluminum turns to be between aluminum and silicon then down between silicon and germanium, between germanium and arsenic, between arsenic and antimony, between antimony and tellurium, between tellurium and polonium, and between polonium and astatine. This is the line between metal and non-metal elements. Metal elements are to the left and down from the line and non-metal elements are to the right and up from the line. Well, that’s not exactly true. There is a line of non-metal elements, Group 8, or Group 18, or Group 0, whichever way you count them, the noble or inert gases that are really an entire Group and category to themselves. Hydrogen is a unique element, the only member of its own Group and category.
The noble gases, or inert gases, have the following properties: For the most part, they do not make chemical combinations with any elements. There have been some compounds made with the noble gases, but only with difficulty. There are certainly no natural compounds with this group. They are all gases at room temperature. They all have very low boiling and melting points. They all put out a color in the visible wavelengths when a low pressure of the gas is put into a tube and a high voltage current is run through the tube. This type of tube is called a neon light whether the tube has neon in it or not. The inert gases are non-metals because they are not metals, but they are significantly different from the other non-metals. As closely akin as all the noble gases are to each other, they should surely be considered a separate group.
By far the largest category of elements on the Periodic Chart is the metal elements. Metals share a set of properties that are not as universal to them as the inert gases. Metal elements usually have the following properties: They have one, two, or three electrons on the outside electron shell. The outside electrons make it more likely that the metal will lose electrons, making positive ions. The ions of metals are usually plus one, plus two, or plus three in charge. Metals tend to lose electrons to become stable. They will attach to other elements with ionic bonds almost exclusively. When metal atoms are together in a group, there is a swarm of semi-loose electrons around the atoms. These electrons move about freely among the metal atoms making what is called an electron gas. The electron gas accounts for the shininess of metals. When there is a smooth surface on the metal it will reflect electromagnetic waves (to include visible light) in an organized manner. The shininess is also called metallic luster. The same electron gas accounts for the cohesive tendencies of metals. Cohesive means the material clings to itself. This property can be easily seen with mercury. Mercury atoms cling to other mercury atoms or other metal atoms with an incredible tenacity. This same cohesion of metals occurs in the solid state. Silver is very malleable. That means that if you hit it, the material would more likely change shape than shatter. At one time US half dollar coins were made of ninety percent silver. It is illegal to deface money, but school children would take a spoon and beat the sides of the silver half dollars until the edges curled inward. When the center became the right size, it was taken out to make a silver ring beaten to fit your finger. Wire is made by pulling metals through a die. The metal coheres to itself so much that it will reshape itself to the shape of the die as it passes through the hole in the die. This property of being able to be pulled through a die to make wire is called ductility (from the Latin ducere, to pull). The presence of the electron gas makes metals good conductors of electricity. Again due to the cohesive property, metals have high melting and boiling points. Almost all metals are solids at room temperature. Metals are usually good conductors of heat. Active metals react with acids. Some very active metals will react with water. Metal elements tend to be denser than non-metals.
The properties of non-metals are not as universal to them as the metals; there is a great deal of variation among this group. Non-metals have the following properties: Non-metals usually have four, five, six, or seven electrons in the outer shell. When they join with other elements non-metals can either share electrons in a covalent bond or gain electrons to become a negative ion and make an ionic bond. When non-metal elements join by covalent bonds, it is usually to other non-metals. Non-metals can attach together with covalent bonds to make a group of (usually non-metal) elements with a common charge called a radical or polyatomic ion. Elemental non-metals often have a dull appearance. They are more likely to be brittle, or shatter when struck. Although not a constant rule, non- metals tend to have lower melting and boiling points than metals and the solids tend to be less dense. Non-metals are not as cohesive as metals and certainly not ductile. Non-metals are not usually good conductors of heat or electricity. Many non-metals form diatomic or polyatomic molecules with other atoms of the same element. Many non-metals have more than one form of the free element, called allotropes, that appear in different conditions. (The word free here means that the element is unattached to other types of atom, not that it has a monetary value of zero.) (I need a link here to Allotropes that used to be in http://chemtutor.com/ms.htm.)
We have pretended that there is a sharp dividing line between the metals and non-metals. This is not the case. The staircase-shaped line between metals and non-metals has several elements on or near it that have properties somewhere between the two categories. By having three electrons in the outside shell, boron should be a metal element. It is not. Boron is more likely to form covalent bonds like a non-metal than donate electrons like aluminum, the next element down the chart in the same group. Aluminum is definitely a metal in most of its traits, but it has its own idiosyncrasy. Aluminum is amphoteric; it reacts with both acids and bases. Silicon, germanium, arsenic, antimony, and tellurium are on the line between metals and non-metals and exhibit some of the qualities of both. These elements do not really comprise a clear-cut category, but, due to the mix of properties they show, they are often lumped into a classification called semi-metals. Many of the elements on the line are semiconductors of electricity, meaning that they have the ability to conduct electricity somewhere between almost none and full conduction. This property is useful in the electronics industry.
We have failed to include hydrogen in any of the categories, for good reasons. Hydrogen just does not match anything else. More than ninety-nine-point-nine percent of hydrogen is just one proton and one electron. A very small proportion (one atom in several thousand) of hydrogen is deuterium, one proton, one neutron, and one electron. An even smaller portion (one hundred atoms per million billion) of hydrogen is tritium, one proton, two neutrons, and one electron. When a hydrogen atom gains an electron, it becomes a negative ion. The negative hydrogen ion, called hydride ion, can be attached to metals, but it is not seen in nature because it is not stable in water. The positive hydrogen ion is what is responsible for acids. There really is no such thing as a (positive) hydrogen ion. Having only a proton and an electron, hydrogen becomes only a proton if it loses its electron. Loose protons attach themselves to a water molecule to make H3O+ ion, a hydronium ion. This hydronium is the real chemical that produces the properties of acids. Elemental hydrogen is a diatomic gas. Except for having a valence of +1, hydrogen has few other similarities with the Group 1 elements. Hydrogen makes covalent bonds between other hydrogen atoms or other non-metals. See hydrogen in the Elements chapter. (need a link to what used to be http://www.chemtutor.com/elem.htm#hydrogen)
GROUPS OR FAMILIES OF THE PERIODIC TABLE
This section is not intended as an exhaustive study of the groups of the Periodic Table, but a quick-and-dirty overview of the groups as a way to see the organization of the chart. Many texts and charts will label the groups with different names and numbers. Chemtutor will attempt to give some standard numbers and identify the elements in those groups so there is no question about which ones we are describing. It is a good idea to have a copy of the Periodic Table available as you go through this section.
Group I (1) elements, lithium, sodium, potassium, rubidium, cesium, and francium, are also called the alkali metal elements. They are all very soft metals that are not found free in nature because they react with water. In the element form they must be stored under kerosene to keep them from reacting with the humidity in the air. They all have a valence of plus one because they have one and only one electron in the outside shell. All of the alkali metals show a distinctive color when their compounds are put into a flame. Spectroscopy (dividing up the spectrum so you can see the individual frequencies) of the colored light from the flame test shows strong emission lines from the elements. The lightest of them are the least reactive. Activity increases as the element is further down the Periodic Chart. Lithium reacts leisurely with water. Cesium reacts very violently. Very few of the salts of Group 1 elements are not soluble in water. Compounds of the alkali metals sodium and potassium are very common in the earth’s crust. Francium is both rare and radioactive.
Group II (2) elements, beryllium, magnesium, calcium, strontium, barium, and radium, all have two electrons in the outside ring, and so have a valence of two. Also called the alkaline earth metals, Group 2 elements in the free form are slightly soft metals. Magnesium and calcium are common in the earth’s crust.
Group III (or 13) elements, boron, aluminum, gallium, indium, and thallium, are a mixed group. Boron has mostly non- metal properties . Boron will bond covalently by preference. The rest of the group are metals. Aluminum is the only one common in the earth’s crust. Group 3 elements have three electrons in the outer shell, but the larger three elements have valences of both one and three.
Group IV (or 14) elements, carbon, silicon, germanium, tin, and lead, are not a coherent group either. Carbon and silicon bond almost exclusively with four covalent bonds. They both are common in the earth’s crust. Germanium is a rare semi-metal. Tin and lead are definitely metals, even though they have four electrons in the outside shell. Tin and lead have some differences in their properties from metal elements that suggest the short distance from the line between metals and non-metals (semi-metal weirdness). They both have more than one valence and are both somewhat common in the earth’s crust.
Group V (or 15) is also split between metals and non-metals. Nitrogen and phosphorus are very definitely non-metals. Both are common in the earth’s crust. In the rare instances that nitrogen and phosphorus form ions, they form triple negative ions. Nitride (N-3) and phosphide (P-3) ions are unstable in water, and so are not found in nature. All of the Group 5 elements have five electrons in the outer shell. For the smaller elements it is easier to complete the shell to become stable, so they are non-metals. The larger elements in the group, antimony and bismuth, tend to be metals because it is easier for them to donate the five electrons than to attract three more. Arsenic, antimony and bismuth have valences of +3 or +5. Arsenic is very much a semi-metal, but all three of them show some semi-metal weirdness, such as brittleness as a free element.
Group VI (6 or 16) elements, oxygen, sulfur, selenium, and tellurium, have six electrons in the outside shell. We are not concerned with polonium as a Group 6 element. It is too rare, too radioactive, and too dangerous for us to even consider in a basic course. Tellurium is the only element in Group 6 that is a semi-metal. There are positive and negative ions of Tellurium. Oxygen, sulfur, and selenium are true non-metals. They have a valence of negative two as an ion, but they also bond covalently. Oxygen gas makes covalent double-bonded diatomic molecules. Oxygen and sulfur are common elements. Selenium has a property that may be from semi-metal weirdness; it conducts electricity much better when light is shining on it. Selenium is used in photocells for this property.
On some charts you will see hydrogen above fluorine in Group VII (7 or 17). Hydrogen does not belong there any more than it belongs above Group 1. Fluorine, chlorine, bromine, and iodine make up Group 7, the halogens. We can forget about astatine. It is too rare and radioactive to warrant any consideration here. Halogens have a valence of negative one when they make ions because they have seven electrons in the outer shell. They are all diatomic gases as free elements near room temperature. They are choking poisonous gases. Fluorine and chlorine are yellow-green, bromine is reddish, and iodine is purple as a gas. All can be found attached to organic molecules. Chlorine is common in the earth’s crust. Fluorine is the most active of them, and the activity decreases as the size of the halogen increases. All of the halogen element gases are poisonous to human beings as a diatomic gas, some more than others. Fluorine and most fluorides (compounds of fluorine) are poisonous. We actually need a small amount of iodine. If we don't get it, our thyroid glands swell up in a condition called goiter. In the US, a small amount of iodine compound is added to table salt, so called iodized salt, so that we can get enough iodine.
The inert gases or noble gases all have a complete outside shell of electrons. Helium is the only one that has only an ‘s’ subshell filled, having only two electrons in the outer and only shell. All the others, neon, argon, krypton, xenon, and radon, have eight electrons in the outer shell. Since the electron configuration is most stable in this shape, the inert gases do not form natural compounds with other elements. The group is variously numbered as Group VII, 8, 8A, 0, or 18. "Group zero" seems to fit them nicely since it is easy to think of them as having a zero valence, that is no likely charge.
The Transition Elements make up a group between what Chemtutor has labeled Group II and Group III. Transition elements are all metals. Very few of the transition elements have any non-metal properties. Within the transition elements many charts subdivide the elements into groups, but other than three horizontal groups, it is difficult to make meaningful distinctions among them. The horizontal groups are: iron, cobalt, and nickel; ruthenium, rhodium, and palladium; and osmium, iridium, and platinum.
Lanthanides, elements 57 through 70, are also called the rare earth elements. They are all metal elements very similar to each other, but may be divided into a cerium and a yttrium group. They are often found in the same ores with other elements of the group. None are found in any great quantity in the earth’s crust. Of the Actinides, elements 89 through 102, only the first three are naturally occurring, the rest being manufactured elements. Of the three naturally occurring ones, only uranium is likely to be referred to in any way in a basic chemistry course. Elements 103 through 109 have been manufactured, and they have been named by the IUPAC (International Union of Pure and Applied Chemistry), but they are not of much importance to this course.
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