Made for the ThinkQuest Internet Challenge 2001 Competition.

One of the most important tools a chemist has is a periodic table.  If you would like to see the periodic table at any time during your mission, use the link below.  We have provided you with both an online version and a downloadable version.

The Atom

Atoms contain three primary particles: protons, neutrons, and electrons.  These particles are referred to as subatomic particles.  Protons are positively charged, electrons are negative, and neutrons are neutral.  Structurally, an atom is set up with a positively charged nucleus containing protons and neutrons that is surrounded by a space where electrons move around the nucleus because of their attraction to the positive charge.  Although electrons orbit the nucleus, they do not do so the same way planets orbit the sun.  Electrons do not orbit the nucleus in a definite path.  Furthermore, electrons move so fast that it is impossible to determine where an electron is at any given time.

 The diagram on the left shows how the nucleus lies on the center of the atom with electrons moving in the space around it.  However, electrons do not move in a defined path.  Therefore, chemists look at electron orbitals as indistinct clouds as is shown by the picture on the right.

The strength of a proton's positive (+) charge is equal to the strength of an electron's negative (-) charge.  All non-bonded elements have the same number of protons as electrons.  Therefore, all elements in their pure form are electrically neutral.  Yet, a proton has a far greater mass than an electron.  The table below summarizes the properties of each of the subatomic particles.

 Particle Location Charge Mass (grams) Mass (AMU) Proton Nucleus +1.602 x 10-19 1.673 x 10-24 1.0073 » 1 Neutron Nucleus 0 1.675 x 10-24 1.0087 » 1 Electron Around Nucleus -1.602 x 10-19 9.109 x 10-28 .0006 » 0

Notice how the mass of each subatomic particle is given in AMU's (Atomic Mass Units).  The AMU unit was created to provide a more simple way of working with the masses of elements and compounds.  As you will see, the atomic mass unit plays an important role in chemistry.

Electrons

Although the exact location of an electron can never be known for sure, there are regions within the electron cloud where electrons most likely to be found.  The ability to determine probable locations of electrons is described by orbitals.  Orbitals describe the probability of finding electrons in certain regions of an atom.  All orbitals have distinctive shapes and sizes.  Every orbital holds up to two electrons.  One electron in that orbital spins clockwise while the other spins counterclockwise, creating a magnetic field.

There are many different kinds of orbitals found in atoms, each having a different shape.  The different kinds of orbitals have been named with the letters s, p, d, and f.  All s orbitals, for example, have a spherical shape.  The p orbitals are dumbbell shaped.  The shapes become more complex with the d and f sublevels.

All atoms contain principal  energy levels (n).  Each energy level can hold a specific number of electrons.  The innermost level is n = 1; the next one out is n = 2, and so on.  The energy of the electron increases from as n increases from 1 to 2 to 3, etc.  Each energy level is divided into one or more sublevels.  There is an interesting and important pattern to these sublevels:  the quantum number n equals the number of sublevels in that principal  level (n=1 has one sublevel, n=2 has two sublevels, and so forth).  The diagram below shows the energy and sublevels of some principal energy levels.

As you can see, the  second principal energy level (n = 2) contains two sublevels:  2s and 2p.  They are called 2s and 2p because they are the types of orbitals found in that sublevel.  So 2p contains p orbitals found in the second principal energy level.  This concept also applies to all other sublevels.  Just as energy increases form n = 1 to n = 2 to n =3, the size of the sublevels increases from lower to higher energy levels.  Therefore, 1s is smaller than 2s which is smaller than 3s.

The 2p energy level, however, does not just have one p orbital, it has three.  It has one p orbital for each axis (x, y, and z).  Look at the table below to see how many orbitals are found in each principal energy level.

 Principal Energy Level Sublevels Orbitals Maximum Electrons n = 1 n = 2 n = 3 n = 4 1s 2s, 2p 3s, 3p, 3d 4s, 4p, 4d, 4f 1s (one) 2s (one) + 2p (three) 3s (one) + 3p (three) + 3d (five) 4s (one) + 4p (three) + 4d (five) + 4f (seven) 2 8 18 32

You may have noticed a pattern in the number of orbitals in each sublevel.  All s sublevels have 1 orbital, all p sublevels have 3 orbitals, all d sublevels have 5 orbitals, and all f sublevels have seven orbitals.

The Periodic Table

The modern periodic table has 109 elements.  That means that everything around us is made up of only 109 different primary pieces!  Each square on the periodic table represents one distinctive element.  There are several different periodic tables, each presenting different information about a particular element, but there are some topics that are common among periodic tables.  First, all elements have an abbreviated name called a symbol.  Hydrogen's symbol is H; helium's symbol is He, etc.  Some elements, on the other hand, have symbols that are more difficult to remember.  That is because many of the symbols found on the periodic table come from that element's Latin name.  Iron's symbol is Fe from its Latin name ferrum.  Use the link at the top of this page to see the periodic table and learn each element's symbol.  Another topic that is commonplace among periodic tables is an element's atomic number, which is a whole number representing the number of protons in the nucleus of an atom.  Next is the atomic mass.  An element's atomic mass is the sum of the masses of that element's protons, neutrons, and electrons.  However, elements sometimes have what are called isotopes.  An isotope is an atom that has the same number of protons as another atom but that has a different number of neutrons.  Carbon, for example, usually has 6 neutrons but is sometimes found with 8. Therefore, the atomic mass that you will find on the periodic table is the average of the masses of the existing isotopes of an element.

Too see a full periodic table, click on the link at the top of the page.

You may be wondering why the periodic table is shaped the way it is.  The periodic table gets its shape form periodic law.  Periodic law is determined by elements' electron  configuration.  Elements with similar bonding properties are found in the same column, called a family or group.  The rows are called periods.  Elements in a period are arranged in order of increasing atomic number.  Elements with the same bonding properties have the same number of valence electrons (electrons in the outermost energy level).  During a chemical reaction, elements either lose or gain valence electrons.  When they do so, they become ions.  Let's look at family 7A.  All elements in that group have 7 valence electrons (notice 7A = 7 valence electrons).  Remember the maximum number of electrons in each energy level mentioned earlier?  Those determine an atom's number of valence electrons.  Fluorine has an atomic number of 9, so it has 9 protons and 9 electrons in its pure form.  The first energy level (n = 1) holds 2 electrons.  That leaves 7 electrons for the second energy level (n = 2), which can hold up to 8 electrons.  Therefore, fluorine has 7 valence electrons!  Now, let's see if that works for chlorine, another element in family 7A.  Chlorine has an atomic number of 17, so it has 17 protons and 17 electrons.  The first energy level holds 2 of those electrons so we are left with 15 electrons.  The second energy level holds another 8 electrons.  Now, we are left with 7 electrons in the third energy level, which can hold up to 18.  There are 7 valence electrons again!

Metals, Nonmetals, and Metalloids

If you look at the periodic table, you will notice three different sections:  metals, nonmetals, and metalloids.  All elements in each category have a few of the same properties.

Most elements on the periodic table are metals.  Metals are found on the left side of the periodic table.  Metals are elements that typically a have high melting point, are ductile (able to be pulled into fine wires), malleable (able to be hammered into thin sheets), shiny, and good conductors of heat and electricity.  Metals tend to lose electrons in a chemical reaction.

Nonmetals are found on the right side of the periodic table.  Nonmetals are elements that have a low melting point, dull surface, break easily, are a poor conductors of heat and electricity, and tend to gain electrons in a chemical reaction.

Metalloids, or semimetals, form a diagonal line on the periodic table between the metals and nonmetals.  Semimetals, as you may have guessed, share properties with both metals and nonmetals but do not accurately fit into either category.

Group Names

In addition to categories such as metals, nonmetals, and metalloids, some families on the periodic table have group names.  The elements in group 1A are alkali metals; the ones in group 2A are called alkaline earth metals.  The elements in group 7A are called halogens and those in 8A are noble gases.  Noble gases have a full outer energy level and tend not to form bonds with other elements.

Periodic Trends

Certain trends, or patterns, exist on the periodic table.  One of these periodic trends is atomic radius.  Atomic radius measures the distance between the center of the nucleus of an atom and the outermost electrons.  Since the exact location of the outermost electrons is not known, this measurement is not precise.  The trend that it follows is that atomic radius increases as you go more toward the bottom-left of the periodic table.  That is because as you move down the periodic table, more energy levels are added, making the element larger.  However, as you proceed to the right side of the periodic table, the radius decreases even though more electrons are added.  The reason is that as you add more electrons to the same energy level, there is a greater force of attraction between the negatively charged electrons and the positively charged nucleus, drawing them closer together.  So fluorine would have a very small atomic radius while cesium would have a relatively large one.

Another periodic trend is found in the elements' ionization energies.  Ionization energy is the energy required to remove the most loosely held electron from an atom.  Metals, as you learned earlier, tend to lose electrons in a chemical reaction.  Thus, they have a far lower ionization energy than nonmetals.  Ionization energy increases as you proceed to the top-right of the periodic table.

Next is electron affinity--the energy change that occurs when an atom gains an electron.  The energy change that occurs becomes progressively more negative toward the top-right of the periodic table, excluding the noble gases.

Finally, there is electronegativity, which is the property of an element that indicates how strongly an atom of that element attracts electrons to itself in a chemical bond.  Elements become gradually more electronegative toward the top-right of the periodic table, excluding the noble gases, with fluorine being the most electronegative.

Nomenclature

Nomenclature is simply the naming of compounds.  This is done by looking at the atoms and their bonds in a compound.  Although you will learn more about bonding later, it is good to know nomenclature in advance.  You will also learn more about how to identify ionic and covalent bonds in Case 3.

Ionic Compound Nomenclature:

An ionic compound is a chemical bond resulting from the transfer of electrons from one bonding atom to another.  Most often, ionic bonds are found between a metal and a nonmetal or a metal and a polyatomic ion.  It is useful to have the common polyatomic ions memorized for more rapid recognition.

 NO3- Nitrate ClO4- Perchlorate IO2- Iodite NO2- Nitrite ClO3- Chlorate IO- Hypoiodite SO4-2 Sulfate ClO2- Chlorite MnO4- Permanganate SO3-2 Sulfite ClO- Hypochlorite OH- Hydroxide PO4-3 Phosphate BrO4- Perbromate CO3-2 Carbonate PO3-3 Phosphite BrO3- Bromate C2O4-2 Oxalate AsO4-3 Arsenate BrO2- Bromite C2H3O2- Acetate AsO3-3 Arsenite BrO- Hypobromite CN- Cyanide CrO4-2 Chromate IO4- Periodate NH4+ Ammonium Cr2O7-2 Dichromate IO3- Iodate
 Notice the pattern in the naming system: per- . . . -ate greatest number of oxygen atoms . . . -ate greater . . . -ite smaller hypo- . . . -ite smallest number of oxygen atoms

Ionic compounds are the easiest to name.  If it is a compound with only two different types of atoms, you state the name of the first atom followed by the name of the second atom; but the name of the second atom must have the suffix -ide.  Thus, the compound NaCl is named sodium chloride.  KI is potassium iodide.  The number of atoms in the compound does not change the name for ionic compounds.  Hence, MgCl2 is still magnesium chloride even though there are two chlorine atoms.

If you have a metal bonded to a polyatomic ion, simply state the name of the metal followed by the name of the polyatomic ion.  AgNO3 is silver nitrate.  You can even have two polyatomic ions bonded to one another, as long as the sum of their charges equals zero.  For example, NH4C2H3O2 is ammonium acetate.

Some atoms, however, can lose different numbers of electrons in a bond.  Listed below are some common examples of these atoms.

 Copper Cu+1 Cu+2 Gold Au+1 Au+3 Cobalt Co+2 Co+3 Iron Fe+2 Fe+3 Chromium Cr+2 Cr+3 Manganese Mn+2 Mn+4 Mn+7 Tin Sn+2 Sn+4 Mercury Hg2+2 Hg+2 Lead Pb+2 Pb+4

Since these atoms can lose different amounts of electrons, you must indicate how many are being lost.  Chlorine will always want to gain one electron when forming an ionic bond.  Therefore, if you are given the formula CuCl, you would know that copper lost one electron to chlorine.  To show that copper lost one electron, the name of the compound is copper (I) chloride.  It has the same naming process as all other ionic compounds except for the Roman numeral (I) to indicate the loss of one electron.  If you are given the formula CuCl2, you can now determine that the name of that compound is copper (II) chloride since copper is losing two electrons (one to each chlorine atom).  Do not worry if you are a bit perplexed at this point, you will obtain a better understanding of nomenclature when you study Case 3 when we discuss bonding.

Molecular Compound Nomenclature:

When naming molecular compounds, you use some of the same rules as you do for ionic compounds.  The first atom in the compound has its name stated first and the second has its name with the -ide suffix.  Yet, unlike ionic compounds, molecular compounds can have two types of atoms bonded in several ways.  For example, carbon and oxygen can bond to form CO (carbon monoxide) or CO2 (carbon dioxide).  Prefixes are used to indicate the number of atoms in a molecule.  Listed below are the common prefixes used to name molecules.

 mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca- 1 2 3 4 5 6 7 8 9 10

When naming molecular compounds, you use a prefix before each of the atoms in the compound.  P2O5, for instance, is called diphosphorus pentoxide.  There is an exception to this rule.  If there is only one of the first atom in the formula, you do not use the prefix mono-, you just write the name of the atom itself.  That is why CO2 is called carbon dioxide not monocarbon dioxide.  Also, if you have a molecule such as H2 or O2, you do not use the di- prefix; you would simply call them hydrogen and oxygen.

Naming Acids:

Acids also have their own set of rules for naming.  If you have a binary acid (an acid made up of two atoms), such as HCl, you take the name of the second atom in the formula and add the suffix -ic. Then, you add the prefix hydro- (from hydrogen) to the beginning.  Finally, you add the word "acid" to the end.  Accordingly, the name of HCl is hydrochloric acid.

If you have a ternary acid (an acid that contains a polyatomic ion), you do not use the prefix hydro-. Additionally,  the suffix -ic is not used in all cases.  What you do is look at the name of the polyatomic ion--if the ion has the suffix -ate, change it to -ic; if the ion has the suffix -ite, it becomes -ous.  Other than that, you keep the name of the ion the same and add the word "acid" to the end.  Here are a few examples to show you how this works.

 HClO4 Perchloric acid HClO3 Chloric acid HClO2 Chlorous acid HClO Hypochlorous acid

The Mole

A mole (abbreviated mol) is defined as the quantity of a substance that has a mass in grams numerically equal to its mass in AMU's.  Interestingly, there are always 6.02 x 1023 atoms or molecules in one mole of that substance.  Consequently, if you have 6.02 x 1023 atoms of carbon, it would have a mass of 12.011 grams because the atomic mass of carbon is 12.011 AMU's.  In addition, if you had one mole of NaCl, it would have a mass of 58.443 grams because the combined atomic masses of Na (22.990) and Cl (35.453) equal 58.443.  The number 6.02 x 1023 is called Avogadro's number in honor of the Italian chemist and physicist Amadeo Avogadro.

 Case 1:  Intro to Chemistry |  Case 2:  Atomic Structure and the Periodic Table |  Case 3:  Bonding  | Case 4:  Intermolecular Forces and Molecular Geometry  | Case 5:  Acids and Bases  | Case 6:  Solutions |  Case 7:  Predicting Products |  Case 8:  Stoicheometry  | Case 9:  Equilibrium  | Case 10:  Nuclear Chemistry Bibliography (AOL users will need to first open Internet Explorer or Netscape Navigator in a new window to use Fullscreen Mode.)