1 The Periodic Table Scientists had identified certain substances as elements and so there were many attempts to arrange the known elements so that there were some correlations between their known properties. Dimitri Mendeleev in 1869 had the idea of arranging elements in order of increasing atomic mass, and, found that elements with similar chemical and physical properties occured periodically. In 1914, Henry Moseley determined that a better arrangement was in order of increasing atomic number. The periodic table is an arrangement of elements in order of increasing atomic number placing those with similar chemical and physical properties in columns.
2 Features of the Periodic Table Groups Vertical columns are called groups. Elements within a group have similar chemical and physical properties. Groups are designated at the top by the numbers 1-8 and by the letters A and B. (Note: group labeling is somewhat arbitrary, so watch out for other designations, particularly with A and B.) A group elements- Representative or main group elements B group elements- Transition elements In addition to the number-letter designation, some groups have their own name. 1A -----> alkali metals 2A -----> alkaline earths 7A -----> halogens 8A -----> noble gases or rare gases Periods Horizontal rows are called periods. Periods are designated by the numbers on the left in the periodic table. The two long rows placed just below the main body of the table are the inner transition elements. Elements 58-71 are the Lanthanide Series
3 The Three Categories of Elements There are three broad categories of elements called (1.) Metals (2.) Non-metals (3.) Metalloids To separate the metals and non-metals we draw a stairstep line to the left of and below B, Si, As, Te, and At. This classification or group is useful because certain properties are associated with each category. Metals -solids at room temperature (except Hg) -metallic luster -malleable and ductile -good conductors of heat and electricity
4 Non-metals -gases or solids at room temperature (except Br 2 ) -variety of color and appearance -brittle solids -insulators (poor conductors) Metalloids -intermediate in properties between metals and non-metals -solids at room temperature -many have more that one structure (one metallic, the other non-metallic) -some are semi-conductors Ions and Molecules Earlier we learned that the atom was comprised of a very small positively charged core of protons and neutrons surrounded by a much large "cloud" of orbiting electrons. While chemistry doesn't usually involve changes in the number of protons or neutrons in an atom, changes in the number of electrons in an atom is central to the science of chemistry.
5 If electrons are removed or added to a neutral atom, a charged particle called an ion is formed. There are two types of ions: Cation - a positively charged ion Anion - a negatively charged ion For example, a neutral sodium atom has a nuclear charge of +11 and contains 11 electrons. If we strip off one electron we form a cation: This process can also be represented in short-hand notation. Let's look at some more examples. A neutral chlorine atom has a nuclear charge of +17 and contains 17 electrons. If we add one electron we form an anion:
6 Another example. Zn likes to lose 2 electrons to make a divalent cation: You can use the periodic table to predict how many electrons an element will lose or gain when it becomes an ion. For example, here are the most stable ionic charges on monoatomic ions: Cl + 1 e - Cl - Cu - 2 e - Cu 2+ N + 3 e - N 3 - O + 2 e - O 2 - F + 1 e - F - Ar + 0 e - Ar 0 K - 1 e - K 1+ Also remember that Aluminum likes to be the cation Al 3+ and Zinc likes to be the cation Zn 2+. Here is a "rough rule" you can use to figure out how many electrons an element will gain or lose: Elements tend to gain or lose electrons to achieve the same number of electrons as the nearest noble gas. For groups in the middle of the periodic table, it is not as simple. After you learn about quantum mechanics, however, you will have a better idea of how to predict the stable ion charges for these groups.
. Chemical Bonds 7 Chemical bonds are the forces that hold atoms together in compounds. They are formed because atoms are not happy with the number of electrons that they have. Only the noble gases (column 8A) are content with the number of electrons. They have the optimum number of electrons and don't like to form chemical bonds. The desire of atoms to gain or lose electrons to get a noble gas number of electrons is what leads to chemical bonding. (It is quite common for chemists to personify the atoms and molecule with which they work! Saying an atom wants an another electron is akin to saying a ball wants to roll down a hill. Later will we examine the energetic and thermodynamic bases of this personification.) There are three types of chemical bonds: Covalent Bonds - atoms are held together by sharing electrons Electrostatic (Ionic) Bonds - cations and anions are held together by electrostatic attractions Metallic Bonds - occurs in metals (similar to covalent bonds) Later, you will be able to determine just how ionic or covalent a bond will be, but for now here are some guidelines to follow: Bonds amongst non-metal atoms are covalent. (For example, a P-S bond is a covalent bond.) Bonds between a non-metal and a metal are ionic (For example, a Na + Cl - bond is ionic) Bonds amongst metal atoms are metallic Metalloid--Non-metal bonds are usually covalent Metalloid--Metal bonds are usually ionic
8 Covalent Bonding In covalent bonding atoms share electrons. Take for example the H 2 molecule. Each hydrogen atom says, "I only need one more electron to be like a noble gas (helium)." Since each hydrogen has only one electron, when two hydrogens get together they can share their electrons. So each hydrogen atom now sees 2 electrons when it is covalently bonded to another hydrogen atom. Pure hydrogen exists as H 2 molecules. The same is true for all of the halogens in column 7A: Pure chlorine exists as Cl 2 Pure bromine exists as Br 2 Pure iodine exists as I 2 Chemists often use the symbol "-" to represent a bond. For example, H-H is a "hydrogen molecule" and Cl-Cl is a "chlorine molecule." The line in between the two atoms means that they are sharing two electrons between them. Let's take oxygen as another example. Oxygen atoms like to combine to form O 2. In this case, each oxygen atom wants 2 more electrons, so when the two oxygen atoms get together they share a total of 4 electrons. We write O 2 as:
9 Chemists call this a double bond. By forming a double bond between them, each oxygen atom can then see as many electrons as a Ne atom has. Now let's look at nitrogen. It also likes to combine to form a diatomic molecule, in this case N 2. Each nitrogen atom, however, wants 3 electrons, so two nitrogen atoms share a total of 6 electrons. We call this a triple bond. You can form molecules from more than one type of atom. Let's look at water. H 2 O consists of two hydrogen atoms sharing their electrons with one oxygen atom. Another example is hydrogen peroxide, H 2 O 2.
Think about hydrogen peroxide and decide on your own if all of the atoms are happy with the number of electons around them. 10 Here is one final example. Carbon atoms want to share 4 electrons, so it is very happy if it can get together with 4 hydrogens to form methane, CH 4. In this example, carbon is sharing 4 electrons with 4 hydrogens and each hydrogen is sharing one electron with carbon
Structural and Empirical Formulas. Structural Formula 11 To avoid confusion, chemists often write the structural formula when identifying a molecule. The structural formula tells you how many of each type of atom are in a molecule and also how they are connected. For example, here is the structural formula of ethanol. Chemical Formula You will also see the term chemical formula. The chemical formula tells you how many of each type of atom are in a molecule. For example, the chemical formula for ethanol is C H O. 2 6 Notice that this is less information than the structural formula (but more compact). You must be careful not to confuse substances that have the same chemical formula. For example, ethanol and dimethyl ether have the same chemicial formula (i.e. C 2 H 6 O). Their chemical formulas are identical, but their structural formulas and their physiological effects are markedly different.
12 Empirical Formulas An empirical formula( simplest formula ) tells us the simplest whole number ratio of atoms in a molecule. For example, hydrogen peroxide's chemical formula is H 2 O 2, but its empirical formula is HO. The chemical formula for glucose is C 6 H 12 O 6, but its empirical formula is CH 2 O, and its structural formula is
13 Molecular Cations and Anions Molecules can also lose or gain electrons to become cations or anions. For example, the NO 3 molecule will gain an electron to form the nitrate anion. If you count up all of the electrons you'll find that all of the atoms feel like neon. Here is the ammonium ion, an example of a molecular cation. The ammonium ion has given up an electron to become a cation.
14 Ionic Bonds Ionic bonds are generally formed when you bring atoms which really want to lose electrons together with atoms which really want to gain electrons. The Na + cation and the Cl - anion are held together by electrostatic or ionic bonds. There is no sharing in ionic bonding. The anion takes the electron for itself and the cation is happy to get rid of its electron. The ions in ionic compounds are arranged in three-dimensional structures. There are no discrete molecules of NaCl. We can only write an empirical formula of NaCl. Here are some other examples. NH 4 Cl : here NH 4 + here Cl - = cation = anion BaCl 2 : here Ba 2+ = cation here Cl - = anion All substances are electrically neutral. We can use this fact to obtain the chemical formula of an ionic compound. Notice that in Na 2 S, two sodium cations were needed to balance the -2 charge of S 2-, making things electrically neutral.