1. Introduction to Chemistry
2. The Periodic Table
3. Quantum Numbers
4. Electron Configuration
5. Chemical Families
6. Oxidation Numbers
7. Chemical Formulas
8. Chemical Names
9. Formula Mass
10. Percentage Composition
11. Reaction Types
12. Balancing Equations
13. The Mole Concept
14. Solution Concentration
16. Kinetic Theory
17. The Gas Laws
18. Enthalpy & Heat
19. Reaction Rates
20. Acids & Bases
21. pH Scale
23. Net Ionic Equations
24. Redox Reactions
25. Organic Chemistry
26. Nuclear Chemistry
5. Chemical Families
A timeline of chemical element discoveries.
Families of elements on the periodic table are in vertical columns called groups.
Through the years, three different systems have been used to number the groups on a Periodic Table.
Families contain elements with similar characteristics, usually determined by the number of electrons in the outer electron energy level.
Chemists are a traditional bunch, so many Periodic Tables are printed with the old and new system. However, the new IUPAC system is now widely accepted.
- The first two used a combination of Roman numerals and letters.
- The old IUPAC system used the letters to designate the left (A) and right (B) side of the Table.
- The old CAS system used the letters to designate the main group elements (A) and transition elements (B).
- The new IUPAC system uses common Arabic numbers, from 1 to 18, to number the columns from left to right across the Table.
Many "trends" in the physical and chemical properties of the elements are directly located to their family and their location on the Periodic Table.
Most metallic elements have a shiny luster, have high melting points, are good conductors of heat and electricity, are malleable (can be pounded into thin sheets), are ductile (can be drawn into wires).
Metals tend to have low ionization energies and lose electrons relatively easily to form positive ions.
Compounds of metals with nonmetals tend to be ionic substances.
Most metal oxides are basic, exhibiting their basicity in two ways:
- metal oxide + water → metal hydroxide
Na2O(s) + H2O(l) → 2NaOH(aq)
- metal oxide + acid → salt + water
NiO(s) + 2HCl(aq) → NiCl2(aq) + H2O(l)
Nonmetals vary greatly in appearance. Most are not lustrous, have low melting points, are poor conductors of heat and electricity, and not malleable or ductile.
Because of their electron affinities, nonmetals tend to gain electrons relatively easily to form negative ions.
Compounds composed entirely of nonmetals are molecular substances.
Most nonmetallic oxides are acidic, exibiting their acidity in two ways:
- nonmetal oxide + water → acid
CO2(g) + H2O(l) → H2CO3(aq)
- nonmetal oxide + base → salt + water
CO2(g) + 2NaOH(aq) → Na2CO3(aq) + H2O(l)
Metalloids touch the metal/nonmetal line on the Periodic Table and have properties between those of metals and nonmetals. They may have some characteristic metallic properties but lack others.
Several metalloids, most notably silicon and germanium, are electrical semiconductors and the principal elements used to manufacture integrated circuits and computer chips.
The 9 Families and Their Characteristics
Group IA (1): The Alkali Metals
The alkali metals have 1 valence electron. They are soft metallic solids, with characteristic metal properties of shiny luster and high thermal and electrical conductivities. The alkali metals have low densities and melting points, increasing and decreasing respectively with increasing atomic mass.
Group IIA (2): The Alkaline Earth Metals
The alkali metals react vigorously with water, producing hydrogen gas and a solution of an alkali metal hydroxide. These reactions are very exothermic, in many cases generating enough heat to ignite the H2, producing a fire or sometimes even an explosion.
Sodium and potassium are abundant in all biological systems.
Hydrogen is in the Alkali Metals Family - but;
Although hydrogen has 1 valence electron and is found on the periodic table in Group 1A, it does not truly belong to any particular group. Because it has no core of electrons producing nuclear shielding, its ionization energy is markedly higher than that of the alkali metals. Where alkali metals lose their valence electron, hydrogen normally shares its valence electron. The fact that hydrogen can actually gain an electron when forming hydrides further illustrates that it is not truly a member of the alkali metal family.
The alkaline earth metals have 2 valence electrons. They have characteristic metal properties but are harder, more dense, and melt at higher temperatures than the alkali metals.
Group IIIA (13): The Boron Group (Earth metals)
Their reaction with water increases down the group. Beryllium does not react with water, even when heated red-hot. Magnesium does not react with liquid water but does react with steam. Calcium and the elements below it react with water at room temperature, although more slowly than the alkali metal adjacent to them on the periodic table.
Magnesium and calcium are essential for living organisms. 99% of the calcium in the human body is found in the skeletal system.
Members of this group have 3 valence electrons. Aluminum, the most familiar and useful member, is found in nature as bauxite - Al2O3.
Group IVA (14): The Carbon Group (Tetrels)
Members of this group have 4 valence electrons. Carbon is unique among the elements because of its ability to combine with itself and other elements to form an almost limitless number of compounds.
Group VA (15): The Nitrogen Group (Pnictogens)
Carbon is found in nature in the form of coal and diamonds. The chemical difference between them is their crystalline structure, which gives them very different physical properties. Diamond is the hardest of all known substances.
Members of this group have 5 valence electrons. Nitrogen makes up about 80% of the air. At normal temperatures and atmospheric pressures, molecular nitrogen is almost inert.
Group VIA (16): The Oxygen Group (Chalcogens)
Commercial production of nitric acid is accomplished using the Ostwald process. Most of the nitric acid produced in the United States is made into fertilizers such as ammonium nitrate. Nitric acid is also used in the manufacture of drugs, dyes, plastics, and explosives. Nitroglycerin and trinitrotoluene (TNT) are two commercial explosives.
Members of this group have 6 valence electrons. There is a change from nonmetallic to metallic character as we move down this group. Oxygen, sulfur, and selenium are typical nonmetals. Tellurium has some metallic properties and is classified as a metalloid. Polonium, which is radioactive and quite rare, is a metal.
Group VIIA (17): The Halogens
The halogens have 7 valence electrons. They are typical nonmetals. Their melting and boiling points increase with increasing atomic number. Fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid.
Group VIIIA (18): The Noble Gases
Halogens are highly active with high electron affinities. In fact, the letter X is sometimes used in chemical equations to indicate any one of the halogen elements. Fluorine gas is so reactive that it is difficult and dangerous to use, requiring specialized lab equipment. Unlike fluorine, chlorine reacts slowly with water to form relatively stable aqueous solutions. Chlorine is often added to drinking water and swimming pools, where the HOCl(aq) that is produced serves as a disinfectant. Halides are all very soluble in water and dissolve to form the hydrohalic acids.
The noble gases have 8 valence electrons. Because the noble gases possess such stable electron configurations, they are exceptionally unreactive. In fact, until the early 1960s the elements were called the inert gases because they were thought to be incapable of forming chemical compounds. Xenon is now known to react with F2 (g) to form the molecular compounds XeF2,
XeF4, XeF6. Only one stable compound of krypton is known, KrF2.
The B Groups (3-12): The Transition Metals
All d and f block metals have 2 valence electrons. Since they are large atoms, the "shelding effect" causes them to have a wide range of possible oxidation numbers when forming compounds.
The noble metals are in this family.
Some resources will identify 10, or even 11 families of elements. This class will recognize only 9.
The difference involves the placement of the Lanthanoid Series (elements 57 - 71) and Actinoid Series (elements 89 - 103). Some references take them out of the transition metals and consider them seperate families. Others place them in a single family known as the Rare-earth Metals.
Size of Atoms
Within each group, atomic radius tends to increase from top to bottom. The major factor influencing this trend is an increase in the principal quantum number, n, indicating the addition of an energy level.
Within each period, atomic radius tends to decrease from left to right. The major factor influencing this trend is an increase in the effective nuclear charge, Zeff. The increasing charge draws the valence electrons closer to the nucleus, decreasing the atomic radius.
There is little change in radii across a d-block period. The size is determined by the s electrons. The attraction produced by the additional protons in the nucleus are, more or less, offset by electron screening.
Removing electrons from a neutral atom to form a cation removes the outer orbitals and decreases the number of electron-electron repulsions, causing the remaining electrons to move closer to the nucleus.
Cations are smaller than their parent atom.
Adding electrons to a neutral atom to form an anion increases the electron-electron repulsions, causing the electrons to move farther from the nucleus.
Anions are larger than their parent atom.
Ionization energy is defined as the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion.
The first ionization energy, I1, is the energy needed to remove the first electron from a neutral atom.
Na(g) → Na+(g) + e−
The second ionization energy, I2, is the energy needed to remove the second electron from a neutral atom.
Na+(g) → Na+2(g) + e−
The greater the ionization energy, the more difficult it is to remove an electron.
I1 is less than I2, which is less than I3 - and so forth. The closer an electron is to the nucleus, the more energy required to remove it from the atom.
Periodic Trends in First Ionization Energies
- Within each period, I1 generally increases with increasing atomic number. The alkali metals have the lowest and the noble gases have the highest.
- Within each group, I1 generally decreases with increasing atomic number.
- The "main-group" elements have a larger range of I1 than the transition elements. Generally, the I1 of the d-block and f-block metals slowly increases from left to right across a period.
- There is a sharp decrease in I1 from the noble gas at the end of one period and the alkali metal at the beginning of the next period (from helium to lithium or from neon to sodium). This supports the idea that only the outermost (valence) electrons are involved in the sharing and transfer of electrons in chemical bonding.
- There is a decrease in I1 when moving from the s subshell to the p subshell in each period (from beryllium to boron or from magnesium to aluminum). The p subshell is farther from the nucleus, decreasing the effective nuclear charge and requiring less energy to remove the electron.
In addition, an atom with a "full" s subshell has some stability, increasing its I1.
- There is a decrease in I1 when moving from an atom with three p electrons to an atom with four p electrons (from nitrogen to oxygen or from phosphorus to sulfur). As stated by Hund's rule, the first three p electrons occupy orbitals singly, thus minimizing electron-electron repulsion. The fourth p electron will require less energy to remove.
Electron affinity is defined as the energy change that occurs when an electron is added to a gaseous atom. This is essentially the opposite of ionization energy.
Ionization energy measures the ease with which an atom loses an electron, where electron affinity measures the ease with which an atom gains an electron.
Cl(g) → Cl+(g) + e−
(ΔE = 1251 kJ/mole)
Cl(g) + e− → Cl−(g)
(ΔE = −349 kJ/mole)
The greater the attraction between an atom and an added electron, the more negative the atom's electron affinity. If an atom were to have a positive electron affinity, an electron would not attach itself to the atom. The nobel gases have positive electron affinities.