Periodic trends are specific patterns found in the periodic table that depict several features of a particular element. Significant periodic patterns include electronegativity, ionization energy, electron affinity, atomic radius, melting temperature, and metallic nature. Periodic trends, which arise from the periodic table’s arrangement, offer chemists with an invaluable tool for swiftly predicting an element’s attributes. These patterns are due to the periodic nature of the elements as well as their similar atomic structures within their various group families.
What are Periodic Trends?
Periodic trends are patterns that can be seen among the chemical elements in the periodic table. These patterns indicate variations in elemental properties as well as atomic structure, including changes in size and radius.
In the periodic table, these modifications take place within each period from left to right and in groups from top to bottom. This is because some characteristics within a time period or group have structural similarities that allow for these tendencies to emerge. These trends are easily mapped to the periodic table, but there are many trends that have exceptions.
Electronegativity Periodic Trends
Electronegativity is a measure of attraction an atom has for bonding electrons. This property reflects how easily an atom can form a chemical bond.
- Electronegativity rises from left to right across a period: It is simpler to remove an electron into the valence shell than to contribute one if it is more than halfway full.
- The electronegativity of a group decreases from top to bottom: This is due to the fact that as atomic number decreases within a group, the atomic radius—the separation between the valence electrons and nucleus—increases.
- The noble gases, lanthanides, and actinides are significant exceptions to the aforementioned trend. The noble gases have a full valence shell and often do not draw electrons. The chemistry of the lanthanides and actinides is more complex and doesn’t typically follow any patterns. As a result, the electronegativity values of noble gases, lanthanides, and actinides are zero.
- Although the transition metals have electronegativity values, there is little variation in their values across periods and within groups. This is because their inability to attract electrons as readily as other elements is impacted by their metallic characteristics.
Ionization Energies Periodic Trends
This is the force required to remove an electron from an atom when it is gaseous, or the propensity of an atom to lose electrons. Conceptually, this is the exact opposite of electron affinity. Each electron in an atom has an associated ionization energy that can be measured. The amount of energy necessary for the removal of the first electron is known as the initial ionization energy. The amount necessary to remove the second electron is second ionization energy, and so forth. Electron shielding is a component that impacts ionization energy. This is the inner electrons’ shielding of the valence electrons from the nucleus.
- Within a period, the ionization energy of the elements increases from left to right. This is because of the stability of the valence shell.
- In general, the ionization energy of elements within a group decreases from top to bottom. This is because electrons are shielded.
- Because of their entire valence shells, the noble gases have extremely high ionization energies.
- It is important to note that helium has the highest ionization energy of any element.
Electron Affinity Periodic Trends
An atom’s electron affinity is its ability to accept an electron. This is a quantitative measure of the energy shift that occurs when an electron is introduced to a neutral gas atom. Like ionization energy, an atom’s electron affinity is measured for each additional electron added.
- Within a given time period, electron affinity shifts from left to right. This is due to a decrease in atomic radius.
- Within a group, electron affinity decreases from top to bottom. The rise in atomic radius is responsible for this.
- The noble gases have positive initial electron affinities, which means they require energy to receive an electron.
Atomic Radius Periodic Trends
Atomic radius is the distance between the nucleus of an atom and its outermost electron shell. This distance is affected by several parameters, including the number of elements and the number of electron shells.
- Across a period of elements, atomic size decreases gradually from left to right. This is because all electrons within a period or family of elements are added to the same shell. However, protons are being added to the nucleus at the same time, making it more positively charged. Because the effect of increasing proton number is greater than that of increasing electron number, there is more nuclear attraction. This means that the nucleus pulls electrons more strongly, drawing the atom’s shell closer to it. Valence electrons are pushed closer to the atom’s nucleus. The atomic radius shrinks as a result.
- Down the group, atomic radius increases. Valance electrons are introduced to shells that are further from the nucleus. As shells are added, electron shielding protects the outer electrons from nuclear attraction from the inner electrons.
Metallic and Non-metallic Properties Periodic Trends
When an element undergoes a chemical reaction, its low ionization energy allows it to lose an electron, exhibiting its metallic nature. On the contrary, the non-metallic property is the ability to gain an electron during a process.
- The elements on the left of the periodic table, in general, exhibit the metallic quality. These are referred to as alkali and alkaline earth metals. The metallic character diminishes from left to right across the periods. As previously stated, the elements on the periodic table’s bottom left have the lowest ionization energies. As a result, they are more reactive than other elements in their groupings. They have the lowest electron affinity as well. As a result, they are the most metallic.
- The metallic character increases from top to bottom across the groups.
Exception: The nonmetal hydrogen (H) is an exception to this rule.
Why is the periodic table arranged the way it is? There are specific reasons, you know. Because of the way we organize the elements, there are special patterns that emerge. Watch the video to have a quick visual recap.
- Pinto, Gabriel. “Using Balls of Different Sports To Model the Variation of Atomic Sizes.” J. Chem. Educ.1998 75 725.
- Qureshi, Pushkin M.; Kamoonpuri, S. Iqbal M. “Ion solvation: The ionic radii problem.” J. Chem. Educ.1991, 68, 109.
- Russo, Steve, and Mike Silver. Introductory Chemistry. San Francisco: Pearson, 2007.
- Petrucci, Ralph H, et al. General Chemistry: Principles and Modern Applications. 9th Ed. New Jersey: Pearson, 2007.
- Atkins, Peter et. al, Physical Chemistry, 7th Edition, 2002, W.H Freeman and Company, New York, pg. 390.
- Alberty, Robert A. et. al, Physical Chemistry, 3rd Edition, 2001, John Wiley & Sons, Inc, pg. 380.
- Kots, John C. et. al, Chemistry & Chemical Reactivity, 5th Edition, 2003, Thomson Learning Inc, pg. 305-309.