Periodic trends

Last updated March 16, 2024

In chemistry, periodic trends are specific patterns that are present in the periodic table that illustrate different aspects of certain elements when grouped by period and/or group. They were discovered by the Russian chemist Dmitri Mendeleev in 1863. Major periodic trends include atomic radius, ionization energy, electron affinity, electronegativity, valency and metallic character. These trends exist because of the similar electron configurations of the elements within their respective groups or periods; they reflect the periodic nature of the elements. These trends give a qualitative assessment of the properties of each element.^[1]^[2]

Summary of trends

Periodic property	Across the period	Down the group
Atomic radius	Decreases	Increases
Ionization energy	Increases	Decreases
Electron affinity	Increases	Decreases
Electronegativity	Increases	Decreases
Valency	Increases	Constant
Nonmetallic character	Increases	Decreases
Metallic character	Decreases	Increases
Nuclear Charge	Increases	Increases

Atomic radius

The atomic radius is the distance from the atomic nucleus to the outermost electron orbital in an atom. In general, the atomic radius decreases as we move from left to right in a period, and it increases when we go down a group. This is because in periods, the valence electrons are in the same outermost shell. The atomic number increases within the same period while moving from left to right, which in turn increases the effective nuclear charge. The increase in attractive forces reduces the atomic radius of elements. When we move down the group, the atomic radius increases due to the addition of a new shell.^[3]^[4]

Ionization energy

The ionization energy is the minimum amount of energy that an electron in a gaseous atom or ion has to absorb to come out of the influence of the attracting force of the nucleus. It is also referred to as ionization potential. The first ionization energy is the amount of energy that is required to remove the first electron from a neutral atom. The energy needed to remove the second electron from the neutral atom is called the second ionization energy and so on.^[5]

Trend-wise, as one moves from left to right across a period in the modern periodic table, the ionization energy increases as the nuclear charge increases and the atomic size decreases. The decrease in the atomic size results in a more potent force of attraction between the electrons and the nucleus. However, suppose one moves down in a group. In that case, the ionization energy decreases as atomic size increases due to adding a valence shell, thereby diminishing the nucleus's attraction to electrons.^[6]^[7]

Electron affinity

The energy released when an electron is added to a neutral gaseous atom to form an anion is known as electron affinity.^[8] Trend-wise, as one progresses from left to right across a period, the electron affinity will increase as the nuclear charge increases and the atomic size decreases resulting in a more potent force of attraction of the nucleus and the added electron. However, suppose one moves down in a group. In that case, the electron affinity will decrease as atomic size increases due to adding a valence shell, thereby diminishing the nucleus's attraction to electrons. Although it may seem that fluorine should have the greatest electron affinity, its small size generates enough repulsion among the electrons, resulting in chlorine having the highest electron affinity in the halogen family.^[9]

Electronegativity

The tendency of an atom in a molecule to attract the shared pair of electrons towards itself is known as electronegativity. It is a dimensionless property because it is only a tendency.^[10] The most commonly used scale to measure electronegativity was designed by Linus Pauling. The scale has been named the Pauling scale in his honour. According to this scale, fluorine is the most electronegative element, while cesium is the least electronegative element.^[11]

Trend-wise, as one moves from left to right across a period in the modern periodic table, the electronegativity increases as the nuclear charge increases and the atomic size decreases. However, if one moves down in a group, the electronegativity decreases as atomic size increases due to the addition of a valence shell, thereby decreasing the atom's attraction to electrons.^[12]

However, in group XIII (Boron Family), the electronegativity first decreases from boron to aluminium and then increases down the group. It is due to the fact that the atomic size increases as we move down the group, but at the same time the effective nuclear charge increases due to poor shielding of the inner d and f electrons. As a result, the force of attraction of the nucleus for the electrons increases and hence the electronegativity increases from aluminium to thallium.^[13]^[14]

Valency

The valency of an element is the number of electrons that must be lost or gained by an atom to obtain a stable electron configuration. In simple terms, it is the measure of the combining capacity of an element to form chemical compounds. Electrons found in the outermost shell are generally known as valence electrons; the number of valence electrons determines the valency of an atom.^[15]^[16]

Trend-wise, while moving from left to right across a period, the number of valence electrons of elements increases and varies between one and eight. But the valency of elements first increases from 1 to 4, and then it decreases to 0 as we reach the noble gases. However, as we move down in a group, the number of valence electrons generally does not change. Hence, in many cases the elements of a particular group have the same valency. However, this periodic trend is not always followed for heavier elements, especially for the f-block and the transition metals. These elements show variable valency as these elements have a d-orbital as the penultimate orbital and an s-orbital as the outermost orbital. The energies of these (n-1)d and ns orbitals (e.g., 4d and 5s) are relatively close.^[17]^[18]

Metallic and non-metallic properties

Metallic properties generally increase down the groups, as decreasing attraction between the nuclei and outermost electrons causes these electrons to be more loosely bound and thus able to conduct heat and electricity. Across each period, from left to right, the increasing attraction between the nuclei and the outermost electrons causes the metallic character to decrease. In contrast, the nonmetallic character decreases down the groups and increases across the periods.^[19]^[20]

Related Research Articles

The alkali metals consist of the chemical elements lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). Together with hydrogen they constitute group 1, which lies in the s-block of the periodic table. All alkali metals have their outermost electron in an s-orbital: this shared electron configuration results in their having very similar characteristic properties. Indeed, the alkali metals provide the best example of group trends in properties in the periodic table, with elements exhibiting well-characterised homologous behaviour. This family of elements is also known as the lithium family after its leading element.

<span class="mw-page-title-main">Chemical bond</span> Association of atoms to form chemical compounds

A chemical bond is the association of atoms or ions to form molecules, crystals, and other structures. The bond may result from the electrostatic force between oppositely charged ions as in ionic bonds or through the sharing of electrons as in covalent bonds, or some combination of these effects. Chemical bonds are described has having different strengths: there are "strong bonds" or "primary bonds" such as covalent, ionic and metallic bonds, and "weak bonds" or "secondary bonds" such as dipole–dipole interactions, the London dispersion force, and hydrogen bonding.

Electronegativity, symbolized as χ, is the tendency for an atom of a given chemical element to attract shared electrons when forming a chemical bond. An atom's electronegativity is affected by both its atomic number and the distance at which its valence electrons reside from the charged nucleus. The higher the associated electronegativity, the more an atom or a substituent group attracts electrons. Electronegativity serves as a simple way to quantitatively estimate the bond energy, and the sign and magnitude of a bond's chemical polarity, which characterizes a bond along the continuous scale from covalent to ionic bonding. The loosely defined term electropositivity is the opposite of electronegativity: it characterizes an element's tendency to donate valence electrons.

The noble gases are the naturally occurring members of group 18 of the periodic table: helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). Under standard conditions, these elements are odorless, colorless, monatomic gases with very low chemical reactivity and cryogenic boiling points.

The periodic table, also known as the periodic table of the elements, is an ordered arrangement of the chemical elements into rows ("periods") and columns ("groups"). It is an icon of chemistry and is widely used in physics and other sciences. It is a depiction of the periodic law, which states that when the elements are arranged in order of their atomic numbers an approximate recurrence of their properties is evident. The table is divided into four roughly rectangular areas called blocks. Elements in the same group tend to show similar chemical characteristics.

The atomic radius of a chemical element is a measure of the size of its atom, usually the mean or typical distance from the center of the nucleus to the outermost isolated electron. Since the boundary is not a well-defined physical entity, there are various non-equivalent definitions of atomic radius. Four widely used definitions of atomic radius are: Van der Waals radius, ionic radius, metallic radius and covalent radius. Typically, because of the difficulty to isolate atoms in order to measure their radii separately, atomic radius is measured in a chemically bonded state; however theoretical calculations are simpler when considering atoms in isolation. The dependencies on environment, probe, and state lead to a multiplicity of definitions.

In physics and chemistry, ionization energy (IE) is the minimum energy required to remove the most loosely bound electron of an isolated gaseous atom, positive ion, or molecule. The first ionization energy is quantitatively expressed as

<span class="mw-page-title-main">Unbinilium</span> Chemical element, symbol Ubn and atomic number 120

Unbinilium, also known as eka-radium or element 120, is a hypothetical chemical element; it has symbol Ubn and atomic number 120. Unbinilium and Ubn are the temporary systematic IUPAC name and symbol, which are used until the element is discovered, confirmed, and a permanent name is decided upon. In the periodic table of the elements, it is expected to be an s-block element, an alkaline earth metal, and the second element in the eighth period. It has attracted attention because of some predictions that it may be in the island of stability.

<span class="mw-page-title-main">Ununennium</span> Chemical element, symbol Uue and atomic number 119

Ununennium, also known as eka-francium or element 119, is a hypothetical chemical element; it has symbol Uue and atomic number 119. Ununennium and Uue are the temporary systematic IUPAC name and symbol respectively, which are used until the element has been discovered, confirmed, and a permanent name is decided upon. In the periodic table of the elements, it is expected to be an s-block element, an alkali metal, and the first element in the eighth period. It is the lightest element that has not yet been synthesized.

The electron affinity (E_ea) of an atom or molecule is defined as the amount of energy released when an electron attaches to a neutral atom or molecule in the gaseous state to form an anion.

<span class="mw-page-title-main">Octet rule</span> Chemical rule of thumb

The octet rule is a chemical rule of thumb that reflects the theory that main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. The rule is especially applicable to carbon, nitrogen, oxygen, and the halogens; although more generally the rule is applicable for the s-block and p-block of the periodic table. Other rules exist for other elements, such as the duplet rule for hydrogen and helium, and the 18-electron rule for transition metals.

In atomic physics, the effective nuclear charge is the actual amount of positive (nuclear) charge experienced by an electron in a multi-electron atom. The term "effective" is used because the shielding effect of negatively charged electrons prevent higher energy electrons from experiencing the full nuclear charge of the nucleus due to the repelling effect of inner layer. The effective nuclear charge experienced by an electron is also called the core charge. It is possible to determine the strength of the nuclear charge by the oxidation number of the atom. Most of the physical and chemical properties of the elements can be explained on the basis of electronic configuration. Consider the behavior of ionization energies in the periodic table. It is known that the magnitude of ionization potential depends upon the following factors:

Size of atom;
The nuclear charge;
The screening effect of the inner shells, and
The extent to which the outermost electron penetrates into the charge cloud set up by the inner lying electron.

<span class="mw-page-title-main">Valence electron</span> An electron in the outer shell of an atoms energy levels

In chemistry and physics, valence electrons are electrons in the outermost shell of an atom, and that can participate in the formation of a chemical bond if the outermost shell is not closed. In a single covalent bond, a shared pair forms with both atoms in the bond each contributing one valence electron.

The inert-pair effect is the tendency of the two electrons in the outermost atomic s-orbital to remain unshared in compounds of post-transition metals. The term inert-pair effect is often used in relation to the increasing stability of oxidation states that are two less than the group valency for the heavier elements of groups 13, 14, 15 and 16. The term "inert pair" was first proposed by Nevil Sidgwick in 1927. The name suggests that the outermost s electron pairs are more tightly bound to the nucleus in these atoms, and therefore more difficult to ionize or share.

The Aufbau principle, also called the Aufbau rule, states that in the ground state of an atom or ion, electrons first fill subshells of the lowest available energy, then fill subshells of higher energy. For example, the 1s subshell is filled before the 2s subshell is occupied. In this way, the electrons of an atom or ion form the most stable electron configuration possible. An example is the configuration 1s² 2s² 2p⁶ 3s² 3p³ for the phosphorus atom, meaning that the 1s subshell has 2 electrons, and so on.

The atomic radius of a chemical element is the distance from the center of the nucleus to the outermost shell of an electron. Since the boundary is not a well-defined physical entity, there are various non-equivalent definitions of atomic radius. Depending on the definition, the term may apply only to isolated atoms, or also to atoms in condensed matter, covalently bound in molecules, or in ionized and excited states; and its value may be obtained through experimental measurements, or computed from theoretical models. Under some definitions, the value of the radius may depend on the atom's state and context.

In chemistry, the valence or valency of an atom is a measure of its combining capacity with other atoms when it forms chemical compounds or molecules. Valence is generally understood to be the number of chemical bonds that each atom of a given chemical element typically forms. Double bonds are considered to be two bonds, triple bonds to be three, quadruple bonds to be four, quintuple bonds to be five and sextuple bonds to be six. In most compounds, the valence of hydrogen is 1, of oxygen is 2, of nitrogen is 3, and of carbon is 4. Valence is not to be confused with the related concepts of the coordination number, the oxidation state, or the number of valence electrons for a given atom.

The lanthanide contraction is the greater-than-expected decrease in atomic radii and ionic radii of the elements in the lanthanide series, from left to right. It is caused by the poor shielding effect of nuclear charge by the 4f electrons along with the expected periodic trend of increasing electronegativity and nuclear charge on moving from left to right. About 10% of the lanthanide contraction has been attributed to relativistic effects.

Core electrons are the electrons in an atom that are not valence electrons and do not participate in chemical bonding. The nucleus and the core electrons of an atom form the atomic core. Core electrons are tightly bound to the nucleus. Therefore, unlike valence electrons, core electrons play a secondary role in chemical bonding and reactions by screening the positive charge of the atomic nucleus from the valence electrons.

The d-block contraction is a term used in chemistry to describe the effect of having full d orbitals on the period 4 elements. The elements in question are gallium, germanium, arsenic, selenium, bromine, and krypton. Their electronic configurations include completely filled d orbitals (d¹⁰). The d-block contraction is best illustrated by comparing some properties of the group 13 elements to highlight the effect on gallium.

References

↑ The Periodic Table I. Structure and Bonding. Vol. 181. 2019. doi:10.1007/978-3-030-40025-5. ISBN 978-3-030-40024-8. S2CID 211038510.
↑ Schrobilgen, Gary J. (2019), Mingos, D. Michael P. (ed.), "Chemistry at the Edge of the Periodic Table: The Importance of Periodic Trends on the Discovery of the Noble Gases and the Development of Noble-Gas Chemistry", The Periodic Table I: Historical Development and Essential Features, Structure and Bonding, Cham: Springer International Publishing, pp. 157–196, doi:10.1007/430_2019_49, ISBN 978-3-030-40025-5, S2CID 213379908 , retrieved 2022-07-02
↑ "atomic and ionic radius". www.chemguide.co.uk. Retrieved 2022-06-30.
↑ Huggins, Maurice L. (1922-04-01). "Atomic Radii. I". Physical Review. 19 (4): 346–353. doi:10.1103/PhysRev.19.346.
↑ "7.4: Ionization Energy". Chemistry LibreTexts. 2014-11-18. Retrieved 2022-07-02.
↑ "Ionization Energy Trend | Science Trends". sciencetrends.com. 2018-05-18. Retrieved 2022-07-02.
↑ Zadeh, Dariush H. (2019-07-26). "Atomic shells according to ionization energies". Journal of Molecular Modeling. 25 (8): 251. doi:10.1007/s00894-019-4112-6. ISSN 0948-5023. PMID 31346734. S2CID 198913558.
↑ Gooch, Jan W., ed. (2007), "Electron affinity", Encyclopedic Dictionary of Polymers, New York, NY: Springer, p. 350, doi:10.1007/978-0-387-30160-0_4245, ISBN 978-0-387-30160-0 , retrieved 2022-07-02
↑ "Electron Affinity Trend | Science Trends". sciencetrends.com. 2018-05-14. Retrieved 2022-07-02.
↑ Chemistry (IUPAC), The International Union of Pure and Applied. "IUPAC - electronegativity (E01990)". goldbook.iupac.org. doi: 10.1351/goldbook.e01990 . Retrieved 2022-06-30.
↑ Bickmore, Barry R.; Wander, Matthew C. F. (2018), "Electronegativity", in White, William M. (ed.), Encyclopedia of Geochemistry: A Comprehensive Reference Source on the Chemistry of the Earth, Cham: Springer International Publishing, pp. 442–444, doi:10.1007/978-3-319-39312-4_222, ISBN 978-3-319-39312-4 , retrieved 2022-06-30
↑ Mullay, John (1987), Sen, Kali Das; Jørgensen, C. K. (eds.), "Estimation of atomic and group electronegativities", Electronegativity, Structure and Bonding, Berlin/Heidelberg: Springer-Verlag, vol. 66, pp. 1–25, doi:10.1007/bfb0029834, ISBN 978-3-540-17740-1 , retrieved 2022-06-30
↑ "21.1: The Elements of Group 13". Libretexts. 2013-11-26. Retrieved 2022-06-30.
↑ Franz, Daniel; Inoue, Shigeyoshi (2016). "Advances in the development of complexes that contain a group 13 element chalcogen multiple bond". Dalton Transactions. 45 (23): 9385–9397. doi: 10.1039/C6DT01413E . ISSN 1477-9226. PMID 27216700.
↑ Manutchehr-Danai, Mohsen, ed. (2009), "valency", Dictionary of Gems and Gemology, Berlin, Heidelberg: Springer, p. 899, doi:10.1007/978-3-540-72816-0_22746, ISBN 978-3-540-72816-0 , retrieved 2022-07-02
↑ Vohr, Hans-Werner, ed. (2016), "Valency", Encyclopedia of Immunotoxicology, Berlin, Heidelberg: Springer, p. 947, doi:10.1007/978-3-642-54596-2_201542, ISBN 978-3-642-54596-2 , retrieved 2022-07-02
↑ Valency. Heidelberg Science Library. 1978. doi:10.1007/978-1-4612-6262-6. ISBN 978-0-387-90268-5.
↑ O’Dwyer, M. F.; Kent, J. E.; Brown, R. D. (1978), O’Dwyer, M. F.; Kent, J. E.; Brown, R. D. (eds.), "Many-electron Atoms", Valency, New York, NY: Springer, pp. 59–86, doi:10.1007/978-1-4612-6262-6_4, ISBN 978-1-4612-6262-6 , retrieved 2022-07-02
↑ Daw, Murray S.; Foiles, Stephen M.; Baskes, Michael I. (1993-03-01). "The embedded-atom method: a review of theory and applications". Materials Science Reports. 9 (7): 251–310. doi: 10.1016/0920-2307(93)90001-U . ISSN 0920-2307.
↑ "C9.1 – Periodic Trends". IGCSE AID. 2018-03-05. Retrieved 2022-07-02.