The Periodic Table
The periodic table of the elements is a systematic arrangement of the chemical elements in a table in order of increasing atomic number. The modern periodic table organises elements so that recurring (periodic) trends in chemical and physical properties are visible. Much of the groundwork for the periodic table was laid by the Russian chemist Dmitri Mendeleev, who in 1869 arranged elements by increasing atomic weight and left gaps where he predicted undiscovered elements would fit; his predictions for several missing elements were later confirmed. The modern periodic table is arranged by atomic number (number of protons).
Definitions and important concepts
- Atomic radius - a measure of the size of an atom. Typical definitions use the distance from the nucleus to the outermost electron shell; in practice atomic radius can be defined in several ways (covalent, metallic, van der Waals radius).
- Ionisation energy - the energy required to remove an electron from an isolated gaseous atom or ion. The first ionisation energy is the energy to remove one electron from a neutral atom in the gas phase. Second, third, etc., ionisation energies refer to removing additional electrons successively and are usually larger than the previous ones.
- Electron affinity - the energy change (often released) when an electron is added to a neutral gaseous atom to form an anion. It indicates how much an element "wants" an extra electron.
- Electronegativity - a dimensionless measure of the ability of an atom (in a molecule) to attract shared electrons. Values commonly referenced (Pauling scale) range from roughly 0.7 (for francium, Fr) to 4.0 (for fluorine, F).
- Group - a vertical column of the periodic table. Elements in the same group have the same number of valence electrons and often similar chemical properties. Groups are numbered 1-18 from left to right.
- Period - a horizontal row of the periodic table. The period number corresponds to the highest principal energy level occupied by electrons in the ground state of the element (for example, elements of period 3 fill the n = 3 energy level).
How position relates to electronic configuration
The periodic table is divided into blocks related to the atomic orbitals being filled: the s-block, p-block, d-block (transition metals) and f-block (lanthanoids and actinoids). An element's position (group and period) helps determine its electronic configuration and valence electrons. For example, phosphorus (P) is in period 3 and group 15 (p-block). Its electron configuration is [Ne] 3s2 3p3, so it has five valence electrons (the 3s and 3p electrons).Trends across a period
As one moves left to right across a given period (constant principal quantum number):
- Atomic radius generally decreases. This is because increasing nuclear charge pulls electrons closer to the nucleus (without a significant increase in shielding within the same shell).
- First ionisation energy generally increases. Atoms more strongly hold their outer electrons when nuclear charge increases and radius decreases.
- Electronegativity generally increases.
- Melting and boiling points often rise to a maximum near the middle of the period (for example around silicon in period 3) and then fall toward the noble gas at the right end, but this pattern depends on bonding types (metallic, covalent, molecular).
- Electrical conductivity increases from the left (metallic conductors) to near aluminium; silicon is a semiconductor and elements to the right are generally poor conductors (insulators/molecular substances).
Period 3: example summary table
| Element (symbol) | Atomic number | Typical chloride(s) | Typical oxide(s) | Outer electrons / Electron configuration (shorthand) |
|---|---|---|---|---|
| Sodium (Na) | 11 | NaCl | Na2O | 1 valence: [Ne] 3s1 |
| Magnesium (Mg) | 12 | MgCl2 | MgO | 2 valence: [Ne] 3s2 |
| Aluminium (Al) | 13 | AlCl3 | Al2O3 | 3 valence: [Ne] 3s2 3p1 |
| Silicon (Si) | 14 | SiCl4 | SiO2 | 4 valence: [Ne] 3s2 3p2 |
| Phosphorus (P) | 15 | PCl3 or PCl5 | P4O6 or P4O10 | 5 valence: [Ne] 3s2 3p3 |
| Sulfur (S) | 16 | SCl2, S2Cl2 (various chlorides) | SO2, SO3 | 6 valence: [Ne] 3s2 3p4 |
| Chlorine (Cl) | 17 | Cl2 (diatomic as an element); forms covalent chlorides with many elements | Cl2O, Cl2O7 (oxides) | 7 valence: [Ne] 3s2 3p5 |
| Argon (Ar) | 18 | - (noble gas, generally forms no stable chlorides) | - (noble gas, generally forms no stable oxides) | 8 valence (filled shell): [Ne] 3s2 3p6 |
Note: Argon is a noble gas with a complete valence shell and is chemically inert under normal conditions.
Trends down a group
As one moves down a group (elements with the same number of valence electrons but increasing principal quantum number):
- Atomic radius increases (electrons occupy higher energy levels further from the nucleus).
- First ionisation energy decreases (outer electrons are further from the nucleus and more easily removed).
- Electronegativity decreases.
- Melting and boiling points often decrease for the alkali metals; other groups show different patterns depending on bonding.
- Density generally increases down many metallic groups (not universal, but commonly observed for alkali metals).
Group 1 (Alkali metals): example summary table
| Element (symbol) | Atomic number | Shorthand electron configuration (valence) | Typical chloride | Typical oxide |
|---|---|---|---|---|
| Lithium (Li) | 3 | [He] 2s1 | LiCl | Li2O |
| Sodium (Na) | 11 | [Ne] 3s1 | NaCl | Na2O |
| Potassium (K) | 19 | [Ar] 4s1 | KCl | K2O |
| Rubidium (Rb) | 37 | [Kr] 5s1 | RbCl | Rb2O |
| Caesium (Cs) | 55 | [Xe] 6s1 | CsCl | Cs2O |
Group 1 elements all form halides (chlorides) with a 1:1 stoichiometry and oxides generally with a 2:1 metal:oxygen ratio (M2O). Down the group: atomic radius increases, first ionisation energy decreases, electronegativity decreases, melting and boiling points generally decrease, and density tends to increase.
Chemical properties of the major groups
Certain groups display characteristic chemical behaviour determined largely by their outer electron configuration. Common group names and notes:
- Group 1 - Alkali metals (one valence electron): very reactive metals, form +1 ions.
- Group 2 - Alkaline earth metals (two valence electrons): reactive metals, form +2 ions.
- Group 13 (three valence electrons): includes aluminium and other elements with mixed metallic/non-metallic behaviour.
- Group 14 (four valence electrons): contains carbon and heavier congeners; shows a range of bonding types.
- Group 15 - Pnictogens (five valence electrons): nitrogen family, includes nitrogen and phosphorus.
- Group 16 - Chalcogens (six valence electrons): contains oxygen, sulfur and heavier congeners.
- Group 17 - Halogens (seven valence electrons): very reactive non-metals, form -1 ions in many compounds.
- Group 18 - Noble gases (filled valence shells): chemically inert under normal conditions.
Note: although hydrogen appears above the alkali metals in many periodic table layouts, hydrogen is not an alkali metal and has unique behaviour.
Applications of periodic trends
Periodic trends allow chemists to predict properties of elements and their compounds: likely oxidation states, sizes, reactivity, types of bonding, typical compounds (oxides, halides), and relative ionisation energies and electronegativities. For example, knowing an element is in group 1 indicates it will typically form a +1 ion and react vigorously with water; knowing an element is in the p-block with a nearly filled shell suggests it may gain electrons to form anions or share electrons covalently.
Activity: Inventing your own periodic table
You have discovered the same set of elements as on Earth but you do not yet have a periodic table. How would you organise the elements so their relationships and trends are clear? Consider which properties you would include (atomic number, atomic radius, ionisation energy, electronegativity, common oxidation states, typical compounds, electron configuration). You might use rows and columns, circular arrangements, or other visual forms. Research alternative periodic-table representations and design one that makes clear patterns for the properties you consider most important. Present your design and reasoning to your class.
Summary
- Elements are arranged in periods (rows) and groups (columns) according to increasing atomic number.
- A group is a column whose elements have similar chemical properties and the same number of valence electrons; a period is a row corresponding to the highest occupied principal energy level.
- Atomic radius is a measure of atom size. Across a period it decreases; down a group it increases.
- First ionisation energy is the energy required to remove one electron from an isolated gaseous atom. Across a period it generally increases; down a group it generally decreases.
- Electronegativity is an atom's tendency to attract electrons; it generally increases across a period and decreases down a group. Fluorine has the highest commonly quoted value (~4.0 on the Pauling scale).
- Groups commonly named: Group 1 - alkali metals; Group 2 - alkaline earth metals; Group 17 - halogens; Group 18 - noble gases. Elements in a group show similar chemical behaviour.
Exercises
Exercise 1
- Use the period 3 summary table and the trends described above to produce a similar table for the elements in period 2.
- Refer to the data table below which gives first ionisation energy (in kJ·mol-1) for elements identified by atomic number Z. Complete the tasks that follow.
Data (atomic number Z - ionisation energy in kJ·mol-1):
| Z | Ionisation energy | Z | Ionisation energy |
|---|---|---|---|
| 1 | 1310 | 10 | 2072 |
| 2 | 2360 | 11 | 494 |
| 3 | 517 | 12 | 734 |
| 4 | 895 | 13 | 575 |
| 5 | 797 | 14 | 783 |
| 6 | 1087 | 15 | 1051 |
| 7 | 1397 | 16 | 994 |
| 8 | 1307 | 17 | 1250 |
| 9 | 1673 | 18 | 1540 |
- a. Fill in the names of the elements for the given atomic numbers.
- b. Draw a graph of atomic number (x-axis) versus ionisation energy (y-axis) and plot the data.
- c. Describe any trends you observe in the plotted data.
- d. Explain why:
- the ionisation energy for Z = 2 is higher than for Z = 1;
- the ionisation energy for Z = 3 is lower than for Z = 2;
- the ionisation energy increases between Z = 5 and Z = 7.
Exercise 2
- Use the Group 1 summary table and the trends described to produce similar summary tables for group 2 and group 17.
- Compare the following elements in terms of the properties listed; explain the differences: 24 12Mg and 40 20Ca.
- Size of atom (atomic radius)
- Electronegativity
- First ionisation energy
- Boiling point
- Study a graph of electronegativity for group 2 elements (Be, Mg, Ca, Sr, Ba) and explain the observed trend.
- Refer to the elements listed below:
- Lithium (Li)
- Chlorine (Cl)
- Magnesium (Mg)
- Neon (Ne)
- Oxygen (O)
- Calcium (Ca)
- Carbon (C)
- belongs to Group 1?
- is a halogen?
- is a noble gas?
- is an alkali metal?
- has an atomic number of 12?
- has four neutrons in the nucleus of its atoms?
- contains electrons in the 4th energy level?
- has all its energy orbitals full?
- will have chemical properties that are most similar (pick a pair and justify)?
End-of-chapter exercises
- For the following statements state whether they are true or false. If false, correct the statement.
- The group 1 elements are sometimes known as the alkali earth metals.
- The group 8 elements are known as the noble gases.
- Group 7 elements are very unreactive.
- The transition elements are found between groups 3 and 4.
- Give one word or term for each of the following:
- The energy that is needed to remove one electron from an atom.
- A horizontal row on the periodic table.
- A very reactive group of elements that is missing just one electron from their outer shells.
- Given 35 80Br and 17 35Cl. Compare these elements in terms of the following properties. Explain the differences in each case.
- Atomic radius
- Electronegativity
