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Chemistry

Chapter 2: 5-9

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WebLecture

An Introduction to the Periodic Table

Outline

The Structure of the Periodic Table

We will return again and again to the periodic table, so get comfortable with it now. This odd layout of information was originally the brainchild of Dmitri Mendelyeev (you'll find this spelled different ways), who decided to arrange elements by two principles:

Mendelyeev didn't have a good reason why the elements should work out this way (we'll spend considerable time discussing the electron orbitals that underlie the actual chemical properties of the elements), but he was able to use his preliminary table to predict the existence of new elements, which were eventually discovered.

Group Group I

Alkaline metals

Group II Alkaline Earth Metals Group II Group IV Group V [Mendeleev's Group VIII
[8A, Transition metals]
Group VI Group VII

Halogens

Noble Gases
Pattern of Combination R2O RO R2O3 RH4 RO2 RH3 R2O5 RO4 RH2 RO3 RH R2O7 Don't combine
Examples Li2O BeO B2O3 CH4

CO2

NH3 FeO4 SiO3 HCl

ClO7

Ar

Chemical Families

A chemical family is a group of elements that combines with some particular element (usually oxygen) in the same way. For example, all alkali metals combine with oxygen in a 2-to-1 ratio, such as Li2O. Because the combinations are similar, we believe that the bonds formed in the combinations are similar and the chemical properties of the elements are also similar with respect to nearly all other elements. In the periodic table, a family is a column designated by a number and letter (1A through 8A, 1B through 8B), and sometimes a distinctive name.

The metals

Metals are shiny, solid, hard, dense, ductile (can be pulled to form wire), malleable (can be hammered into different shapes), and readily conduct heat and electricity. Later on, we'll discover that metal atoms don't hang onto their electrons very well, and the flow of electrons through a sample of metal is responsible for many of its properties.

1A: Alkali Metals This is one end of the table; all members of this family are very reactive. This means that they tend to form chemical bonds with other elements and give off heat energy as they do so. The alkali metals have one loosely held electron in thier outer orbit, and they are willing to trade it in order to form a closer relationship with many other elements. When they combine with halogens (on the other end of the table), they form salts. As mentioned, these bond 2-to-1 with oxygen.

2A: Alkaline Earth Metals These elements are slightly more stable, but like their near-relative Alkali Metals, are reactive enough that they will combine with many different elements, and are almost never found in nature except in compounds. The members of this family bond 1-to-1 with oxygen, so we get molecules like MgO.

Transition metals

1B - 8B These are the metals that are familiar to you, like gold, silver, nickle, copper, zinc, and iron. They all have an electron structure that allows one or more electrons to easily shift orbits within the atom, changing some of its energy characteristics, like the ability to absorb and emit light of particular colors. This makes them important as pigments or coloring agents in paints, as sensors in the cones of your retina, and as energy conversion mechanisms in chlorophyll and beta-carotene, but doesn't affect their chemical properties.

Non-metals

Non-metals are mostly solids or gasses. Solid non-metals are brittle, dont' change shape easily, and are poor conductors of heat or electricity. The metalloids have characteristics of both types: they are somewhat conductive and somewhat ductile.

3A: Boron Family One metalloid, the rest metals.

4A: Carbon Family One non metal (at the top), two metalloids, and the rest metals. This family is unique in having the most forms of the pure element because its elements can combine with up to four other atoms each. As a result, carbon (for example), hast the allotropes or forms graphite, diamonds, and buckyballs. These are all samples of pure carbon, but they have different boiling and melting points, malleability, and densities.

5A: Nitrogen Family A couple of non-metals, a couple of metalloids, and one metal.

6A: Oxygen Family Mostly non-metals, one metalloid, one metal....the last one.

The most reactive family

7A: Halogens These are the most reactive elements in the bunch; they will combine violently with the alkali metals. Later on, when we look at their electron configurations, you'll see why--but for now it is enough to realize that halogens want something alkali metals have, and will combine with the alkalis to get it.

The least reactive family

8A: Noble Gases These elements are self-sufficient isolationists. Under extreme circumstances, they will actually form a kind of bond with other elements (xenon will combine with fluorine), but normally the noble gas atoms do not combine with other atoms, not even each other

Modern Periodic Tables

There are a number of online periodic tables that not only displace some information "up front", but which make it easy to drill down into more detailed information. Some of the best are

  1. Webelements at the University of Sheffeld in England.
  2. The Dynamic Periodic Table, withi similar information.
  3. The Photographic Periodic Table, with pictures of each of element.

There are also several mobile platform (iPhone, Blackberry, etc) applications which you may find useful in studying the periodic table on the go. We will be revisiting different representations of this information when we study electron configurations.

Molecules and Ions

The central study of chemistry is the arrangement of atoms in molecules: what kinds of arrangements can exist, how they can be changed, and how much energy is required to make the change. In this chapter, we look at the types of bonds, or connections, that atoms can have with each other.

Bonds are a result of the electrical charges on the atomic components. As with any kind of electrically charged particles or objects, like charges repel each other and opposite charges attract each other.

The protons in the nucleus are all positively charged. We are not here concerned with how they stick together in the nucleus; the weak and strong nuclear forces and transformations between protons and neutrons are more properly the study of physics. But taken as a whole, the positive charge at the center of the atom holds the atom's electrons in their orbits and repels the positive charges of the nuclei of other atoms.

The electrons on the "outside" of the atom repel each other, and are in constant motion in particular relationships as a result of that repulsion. The outermost of these can be attracted by the positive charge of another atom's nucleus. The interplay of electrical forces between one atom's electrons, its own positive nucleus, the electrons of another atom, and that atom's positive nucleus are responsible for all chemical bonding.

Electron Proton Attraction

Molecules and bonds

Atoms are held together in molecules by four different kinds of bonds. When the total number of electrons equals the total number of protons in the molecule, the molecule is electrically neutral, otherwise, the molecule is an ion. A positively charged ion is called a cation; a negatively charged ion is an anion.

Covalent bonds hold together atoms which share electrons. These are strong bonds, not easily dissolved in solvents like water. Water is held together by covalent bonds, but the hydrogen atoms are on one side of the oxygen atom, so there is a side of the water molecule which is positive, and one which is negative. Such an molecule is called polar, and will move so that its negative side is oriented to the positive side of another polar molecule. [This gives water, in particular, a number of interesting properties, like surface tension].

Sugar is held together by covalent bonds, and is neutral and non-polar, althought temporary concentrations of charge occur as electrons move around in the molecule.

Ionic bonds hold together atoms of opposite electrical charges (one atom has to lose, the other gain, one or more electrons). Table salt (NaCl) is ionic, and dissolves into Na+ and Cl- ions (charged atoms) in water.

Hydrogen bonds are formed by hydrogen atoms in covalent bonds with another atom. Since the hydrogen's single electron is shared with the other atom, its proton's positive charge is unshielded on the open side (opposite the bond). This area of positive charge can attract any polar molecule's negative area. Hydrogen bonds between molecules of water account for water's ability to absorb and hold heat and to adhere to non-water surfaces. Cohesion (attraction between molecules f the same substance) also accounts for water's high surface tension.

Van der Waals forces are similar to hydrogen bonds, but weaker. Since the atom with the positive charge isn't hydrogen and still has some electrons close its nucleus, its positive charge is partially neutralized by its own electrons. You might think of a hydrogen bond as a special case of a van der Waals force, stronger than the others because it doesn't have any electrons shielding its single-proton nucleus.

Bond Strength

All bonds between molecules store energy; therefore all bonds require energy input to make the bond. When bonds are broken, the enrgy is released and can be used to form new bonds, or dispersed as heat, energy, light or sound or kinetic energy (energy of motion).

In general, covalent bonds are the strongest type of bond (a very weak covalent bond can be weaker than a very strong ionic bond). The strength of the bond depends on which electrons are shared, and how many are shared. Single bonds involve each atom sharing one electron (two electrons are shared in each single bond). A double bond involves each atom sharing two electrons (four electrons shared in the bond). Obviously, because it takes energy to pull each electron out of the shared orbital, double bonds can be up to twice as strong as single bonds.

In ionic bonds, the strength of the bond depends on the charges q+ and q-, and the distance D betwen the bonds. Electrical attraction is described by Coulomb's law, where the force of the attraction is

F e   =   k   ( n + e ) ( n e ) d 2 where k is a universal constant, n+ is the number of positive charges on the cation in electron units, n- is the number of negative charges on the anion, e is the charge in coulombs per electron, and d is the distance between the ions (usually determined from the center of the ion).

As the charges increase (for example, from Na+ to Cu3+), the bond becomes stronger. As the distance D increases (and the size of the atoms involved depends on the the number of electrons in the atom), the bond becomes stronger. Obviously, there is a range of strengths possible for ionic bonds.

Hydrogen bonds are uniform in charge strength, since you have one lone proton exerting the electrical field which gives rise to the force, but in practice, the bonds will fary depending on how close the two atoms involved can get--and that will depend on the size and shape of the rest of their molecules.

Van der Waals bonds will depend on the element type (and hence the electron structure) of the atom involved.

One point to keep firmly in mind is that any change in electron configuration, whether it is the ionization of the atom, or its bonding with other atoms in any of the four types of bonds, results in a change in the chemical characteristics of the atom. Since the current electron configuration determines how much energy is required to change to a different configuration, everything from the melting point of the substance to its density and ability to dissolve is affected by that configuration. While some changes are predictable (graphite melts at a lower temperature than diamond, for example, because it involves fewer covalent bonds between all the atoms in the substance), the characteristics of most products must be experimentally determined.

Naming Ionic and Molecular Compounds

Alchemists and early chemists, like early botanists, often had different names for the same substance, or used the same name for more than one substance. This made it difficult to communicate the results of experiments and experiences. During the 19th century, chemists standardized a naming convention for many ions; during the twentieth century, naming conventions for organic compounds were developed. Now we use both symbols and names that uniquely identify molecules by their content and structure.

Different types of formulae are used, depending on the information to be conveyed. Structural formulas show the arrangement of atoms in the molecule. Molecular formulas list the contents of the molecule, listing the positive ion first. Ring structure diagrams are used for carbon rings, like those in glucose. Discovering the actual distribution of the atoms within a molecule is a major part of modern chemical research.

The table below shows the rules for naming ionic compounds. Cations are positively charged ions. Anions are negatively charged. Names ending in -ite and -ite for the same base group means that it combines with different numbers of oxygen atoms. Nitrate is NO3-, nitrite is NO2-.

Type of ion or molecule Naming convention Example Formula Name
Monatomic cation Name of metal element Na+sodium cation or sodium ion
Monatomic cation, higher charge Roman numeral OR stem + ic Fe3+ Fe(III), iron(III) or ferric
Monatomic cation, lower charge Roman number OR stem + ous Fe2+ Fe(II), iron(II) or ferrous
Monatomic anion Stem of element name + ide O-2 oxide
Nonmetal oxoanion, third/fourth form (most oxygen) Per + stem + ate ClO4- perchlorate
Nonmetal oxoanion, higher oxygen Stem + ate ClO3- chlorate
Nonmetal oxoanion, lower oxygen Stem + ite ClO2- chlorite
Nonmetal oxoanion, third/fourth form (least oxygen) Hypo + stem + ate ClO- hypochlorite
Ionic compound Cation name anion name CaS calcium sulfide
Hydrogen containing molecular compound Hydrogen +second element name + ide HF hydrogen fluoride
Binary molecular compound Greek prefix for number +first element name Greek prefix of number +second element name + ide NO3
N2O
N2O3
nitrogen trioxide
dinitrogen oxide
dinitrogen trioxide
High oxygen oxoacids Anion -ate changes to -ic HClO4 perchloric acid
Lower oxygen oxoacids Anion -ite changes to -ous HClO2 chlorous acid
Common names Must be learned CH4
NH3
methane
ammonia

For example, to figure out the formula for a metal hydroxide, look at the charge on the metal. Calcium is a group 2A ion and has 2 electrons in the outer shell to lose, so it forms Ca2+. The oxide ion OH- has a single negative charge. To form a neutral ion with Ca2+, we need two negative charges, so we have to have 2 oxide ions. The resulting molecule has a calcium ion and two hydroxide ions. To show this, we put parentheses around the hydroxide ion: Ca(OH)2.

Practice with the Concepts

Reading the Periodic Table

Use the Periodic Table on the inside front cover of your text, or a web-based Periodic Table such as WebElements. For the element below, what is the

  • natural state at room temperature,
  • the symbol,
  • the type of element metal (non-metal, noble gas, etc.)
  • the atomic mass
  • the number of protons
  • the likely number of neutrons for the most common isotope
Mercury

Element Characteristics

The information on the periodic table is not a complete description. Can you think of other characteristics that might be used to uniquely identify a given element?

Discussion Questions

Optional Readings

Check out the different forms of the period table that have been proposed. Each has some advantages over the standard form in revealing specific relationships.