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Topic 3.2 Periodic Trends - Coggle Diagram
Topic 3.2 Periodic Trends
Essential idea:
Elements show trends in their physical and chemical properties across periods and down groups of the periodic table.
understanding
Atomic radius, ionic radius, ionization energy, electron affinity and electronegativity all show vertical and horizontal trends in the periodic table.
The trends in metallic and non-metallic behaviour are due to the above trends.
Across a period the nature of the oxides changes from basic through amphoteric to acidic.
Apply their knowledge to:
Predict and explain the metallic and non-metallic behaviour of an element based on its position in the periodic table.
Discuss the similarities and differences in the properties of elements in the same group, (limited to the alkali metals, group 1, and the halogens, group 17).
Construct equations to explain the changes in pH for the reactions of sodium oxide, Na2O, magnesium oxide, MgO, phosphorus(V) oxide, P4O10, and the oxides of nitrogen and sulfur with water.
identification
Atomic radius
The sizes of different atoms can be compared using their atomic radii. Section 9 of the IB Chemistry data booklet has data for the atomic and ionic radii of the elements in the periodic table
Ionic radius
Metallic elements tend to lose electrons from their atoms to form positively charged ions (cations). Examples include Na+, Mg2+ and Al3+. Non-metallic elements, on the other hand, tend to gain electrons to form negatively charged ions (anions), such as S2− and Cl−. Ions such as these are described as being isoelectronic.
First ionisation energy
The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. It can be thought of as a measure of the attraction between the positively charged nucleus and the negatively charged outer valence electrons. The general equation for the first ionisation energy of an atom is shown below. Note that the positive sign for the ionisation energy (IE) means that energy is required to remove an electron from the attraction of the nucleus.
Electronegativity
Electronegativity is defined as the attraction of an atom for a bonding pair of electrons. It is measured on the Pauling scale, a relative scale that assigns fluorine a value of 4.0 and francium a value of 0.7. Electronegativity values for the elements can be found in section 8 of the IB Chemistry data booklet.
Electron affinity
The general trend down a group is decreasing electron affinity. The additional electron gained is entering an energy level further from the nucleus. This added electron has a weaker attraction to the nucleus and therefore releases less energy when added. There are subtle variations in the trend of electron affinity, but the overall trend is as described above.
Melting point
The melting point of a substance depends on its structure and bonding. Across a period, the structure and bonding gradually change from metallic to giant covalent to molecular covalent. Note that these types of bonding will be covered in more detail in topic 4. Here, we will only take a brief look at each type of structure and bonding in the elements of period 3 (Na to Cl).
Metals and non-metals
Metals are characterised by having fewer valence electrons, larger atomic radii, lower electronegativity values and lower ionisation energies. They have a tendency to lose their valence electrons relatively easily to form positive ions (cations). Non-metallic elements are characterised by having more valence electrons, smaller atomic radii, higher electronegativity values and higher ionisation energies. They tend to gain electrons to form negative ions (anions) instead of losing electrons to form positive ions.
