Atomic Radius in Periodic Table in Basic Chemistry - Overview, Example, Types & Trends, FAQs

Q: What is the order of atomic radius?

Atomic radius is of the order of 10 −8 cm and nuclear radius is of the order of 10 −13 cm. The fraction of atom occupied by nucleus is: 10 −5 .

Edited By Team Careers360 | Updated on Jul 02, 2025 04:25 PM IST

Atomic size is defined as the distance between an atom's nucleus and its outermost shell. In basic chemistry, it is also denoted by the atomic radius which means the shortest distance between an atom's nucleus and its outermost shell. In a molecule, an atomic radius is half the distance between neighbouring atoms of the same element.

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This Story also Contains

Methods to calculate atomic radius
Atomic radii are classified as
Atomic Radius r=d/2
Size of Hydrogen Atom
Types of Radius in Relation to the Different Types of Bonds
Trends in the Periodic Table

Atomic Radius in Periodic Table in Basic Chemistry - Overview, Example, Types & Trends, FAQs

As the size of an atom is of the order 1.2 ×10^-10 m, measuring the atomic radii of chemical elements is a difficult undertaking. Because the electron cloud that forms an atom's shell has no fixed shape, determining the atomic size of an atom is challenging. As a result, we might conclude that determining the size of each particular atom is essentially impossible.

An element's atomic radius is a measurement of the size of its atoms, which is typically the mean distance between the nucleus centre and the boundary of the electrons' surrounding shells. Due to the lack of a well-defined border, multiple non-equivalent definitions of the atomic radius exist.

Methods to calculate atomic radius

X-ray or other spectroscopic methods are used to determine an atom's atomic radius. The atomic radii of elements vary in a predictable fashion across the periodic table. The nuclear charge and energy level can be used to explain this tendency.

This is comparable to the concept of a circle's radius, in which the nucleus serves as the circle's centre and the electron's outermost orbital serves as the outside border of the circle.

According to the principle, the radius is derived from the distance between the nuclei of two bonded atoms. Atomic radii are thus determined by the bonds that they make. The radius of an atom varies depending on the bond it makes, hence there is no fixed radius for an atom.

Atomic radii are classified as

Ionic radius
Covalent radius
Metallic radius
Van Der Waals radius.

Atomic Radius d/2

Atomic Radius r=d/2

d – Distance between the nuclei of two atoms.

Size of Hydrogen Atom

The van Der Waals radius of the hydrogen atom is 120pm.

Atomic Radius of H2

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Types of Radius in Relation to the Different Types of Bonds

1. Van Der Waals Radius

The Van Der Waals radius is a measurement of the size of an unbound atom, whether chemically, ionically, or covalently. It is defined as half of the distance between two equal and non-bonded atoms at their closest distance.

2. Ionic Radius

The radius of an atom creating an ionic bond or an ion is called the ionic radius.
Because the electrons and nucleus are constrained by atomic bonds, ions and atoms do not have a distinct shape.
The ionic radius is measured in Armstrong (A◦) or Pico meters (pm). The typical radius can be anywhere between 30 and 200 pm.
The ionic radius is not constant; it varies depending on the spin state of the electrons, the coordination number, and a variety of other factors.
The size of an ion increases as the number of coordinates increases. An ion with a high spin state of an electron has a larger ionic size than an ion with a low spin state of an electron.
If the charge of the ion is taken into account, the ion with a positive charge will be smaller than the ion with a negative charge.

3. Metallic Radius

The metallic radius is the distance between two atoms connected by a metallic bond. In a metallic cluster, the metallic radius is half of the total distance between the nuclei of two adjacent atoms.

4. Covalent Radius

An atom's covalent radius is defined as the radius of an atom that is in a covalent bond with another atom(s) of the same element.
By comparing the bond lengths of two covalently bonded atoms of the same kind, the covalent radius of an atom may be calculated: if the two atoms are of the same kind, the covalent radius is just one half of the bond length.
While this is straightforward for some molecules, such as Cl2 and O2, in others, the covalent radius must be calculated by measuring bond distances to atoms whose radii are already known (for example, a C–X bond, where C's radius is known).

Related Topics

Trends in the Periodic Table

There are various trends which can be studied with respect to periods and groups in periodic table.

Across a Period

As we move from left to right in a period, the Covalent and Van der Waals radii fall as the atomic number increases. In a period, the alkali metals at the far left of the periodic chart have the largest size. The halogens on the periodic table's far right have the smallest size. Nitrogen has the smallest atomic size. After nitrogen, the atomic size of oxygen increases, then drops for fluorine. The atoms of inert gases are bigger than those of the halogens before them.
The nuclear charge grows by one unit in each subsequent element as we move from left to right in a period, while the number of shells remains constant. The electrons in all shells are drawn closer to the nucleus by the increased nuclear charge. As a result, each individual shell becomes smaller. As we move from left to right in a period, this causes the atomic radius to decrease.
As we progress from halogens to inert gases, the atomic radius abruptly increases. Because inert gases have entirely filled orbitals, this is the case. As a result, the inter-electronic is at its greatest. Because they do not form covalent bonds, we express atomic size in terms of Van der Waals radius. The radius of Van der Waals is bigger than the radius of covalent bonds. As a result, the atomic size of inert gas in a period is substantially larger than the halogen before it.

Down a Group

The atomic radii of elements rise as the atomic number in a group increases from top to bottom. The principal quantum number increases as we proceed down the group. At each subsequent element, a new energy shell is added. The valence electrons are moving away from the nucleus and away from the nucleus. As a result, the nucleus's attraction to the electron decreases. As a result, the atomic radius increases.

Frequently Asked Questions (FAQs)

1. What is the atomic radius in the periodic table?

The atomic radius is defined as the smallest distance between the nucleus of an atom and its outermost shell.

2. Explain Van der Waals radius?

Van der Waals Radius is half the distance between the nuclei of two similar non-bonded solitary atoms or two near identical atoms belonging to two neighbouring molecules of the same element.

3. Which is the smallest element in the periodic table?

The smallest element in the periodic table is Helium.

4. What is the order of atomic radius?

Atomic radius is of the order of 10⁻⁸ cm and nuclear radius is of the order of 10⁻¹³ cm. The fraction of atom occupied by nucleus is: 10⁻⁵.

5. What is meant by ionic radius?

The ionic radius is the distance between the nucleus and the electron in the outermost shell of an ion.

6. How does the concept of atomic radius apply to allotropes of an element?

Allotropes are different structural forms of the same element. While the atomic radius of the individual atoms remains the same, the arrangement and bonding in different allotropes can lead to variations in interatomic distances. For example, the carbon-carbon distances in diamond, graphite, and fullerenes differ due to their unique structures.

7. What is the relationship between atomic radius and atomic volume?

Atomic volume is proportional to the cube of the atomic radius. As atomic radius increases, atomic volume increases much more rapidly. This relationship is important in understanding trends in properties like density and molar volume across the periodic table.

8. How does atomic radius affect the formation of molecular orbitals?

Atomic radius influences molecular orbital formation by affecting the extent of orbital overlap. Atoms with similar sizes tend to form stronger covalent bonds due to better orbital overlap. When there's a significant size mismatch, orbital overlap may be less effective, potentially leading to weaker or less stable bonds. This concept is important in understanding bonding in diatomic molecules and more complex structures.

9. Why do transition metals often form colored compounds, and how does this relate to atomic radius?

The formation of colored compounds by transition metals is related to d-orbital splitting, which is influenced by atomic radius. The size of the metal ion affects the strength of the crystal field created by surrounding ligands. This, in turn, determines the energy difference between d-orbitals, which corresponds to the wavelength of light absorbed and thus the color observed.

10. What is metallic radius, and how does it differ from other types of atomic radii?

Metallic radius is half the distance between adjacent metal atoms in a crystal lattice of the pure metal. It's typically larger than covalent radius because metallic bonding involves delocalized electrons, allowing atoms to pack more closely. Unlike covalent or van der Waals radii, metallic radius is specific to solid metallic structures.

11. How does the concept of atomic radius apply to non-spherical atoms?

The concept of atomic radius assumes atoms are spherical, which is a simplification. For non-spherical atoms or molecules, different radii may be defined along different axes. In organic chemistry, for instance, van der Waals radii are often given separately for the length and width of elongated atoms or groups.

12. How does atomic radius affect the crystal structure of solids?

Atomic radius plays a crucial role in determining crystal structure. The ratio of cation to anion radii in ionic compounds, for instance, influences the coordination number and the type of crystal structure formed. In metals, atomic radius affects how closely atoms can pack together, influencing properties like density and melting point.

13. What is the significance of atomic radius in understanding the periodic law?

Atomic radius is a key property that demonstrates the periodic law, which states that the properties of elements are periodic functions of their atomic numbers. The periodic trends in atomic radius help explain many other periodic trends in physical and chemical properties, making it fundamental to understanding the organization of the periodic table.

14. Why do atoms of the same element can have different radii in different compounds?

Atoms of the same element can have different radii in different compounds due to variations in oxidation state, coordination number, and the nature of bonding. For example, the radius of an iron atom in Fe2+ is different from Fe3+, and the radius of a carbon atom in a single bond differs from one in a double bond.

15. How is atomic radius measured?

Atomic radius cannot be measured directly for isolated atoms. It is typically inferred from measurements of bond lengths in molecules or from the distances between atoms in crystal structures. Different measurement methods can lead to slightly different values.

16. What is Slater's rule, and how does it relate to atomic radius?

Slater's rule is a method for estimating the effective nuclear charge experienced by electrons in an atom. It takes into account the shielding effect of inner electrons. This rule helps explain atomic radius trends by quantifying how the effective nuclear charge changes across the periodic table, directly influencing the size of atoms.

17. How does the concept of atomic radius apply to elements in extreme conditions, like very high pressures?

Under extreme conditions like very high pressures, atomic radii can change significantly. High pressure can compress electron orbitals, effectively reducing atomic size. This can lead to changes in chemical behavior and even the formation of new phases or compounds that are not stable under normal conditions.

18. How does the concept of atomic radius apply to isotopes of an element?

Isotopes of an element have the same number of protons and electrons but different numbers of neutrons. While the chemical properties and electron configuration remain largely the same, there can be slight differences in atomic radius due to changes in nuclear size and mass. These differences are usually very small and often negligible in chemical contexts.

19. How does the concept of effective nuclear charge relate to atomic radius?

Effective nuclear charge is the net positive charge experienced by valence electrons. As effective nuclear charge increases, it pulls electrons closer to the nucleus, decreasing atomic radius. This explains why atomic radius generally decreases across a period, as the effective nuclear charge increases with each additional proton.

20. What is the significance of atomic radius in nanotechnology?

Atomic radius is crucial in nanotechnology because it determines the size of nanostructures and influences their properties. Understanding atomic radii helps in designing and manipulating materials at the nanoscale, where quantum effects become significant and properties can differ dramatically from bulk materials.

21. How does atomic radius affect the formation of hydrogen bonds?

Atomic radius influences hydrogen bond formation indirectly through its effect on electronegativity. Smaller, more electronegative atoms like fluorine, oxygen, and nitrogen form stronger hydrogen bonds because they create a greater partial negative charge on the bonded hydrogen. This makes the hydrogen more attractive to other electronegative atoms.

22. How does atomic radius affect the strength of chemical bonds?

Generally, smaller atomic radii lead to stronger chemical bonds. This is because the valence electrons of smaller atoms are closer to the nucleus and held more tightly, allowing for greater orbital overlap and stronger bonds. However, bond strength also depends on other factors like electronegativity and the type of bond formed.

23. Why do some elements, like francium, have such large atomic radii?

Francium, the largest alkali metal, has a very large atomic radius due to its position at the bottom of Group 1. It has seven electron shells, with the outermost electron very far from the nucleus. Despite its high nuclear charge, the shielding effect of inner electrons results in a weak attraction for the outermost electron, leading to a large atomic radius.

24. How does atomic radius affect the solubility of elements?

Atomic radius can influence solubility, particularly in ionic compounds. Smaller ions generally have higher charge density, leading to stronger ion-dipole interactions with water molecules and higher solubility. However, solubility is a complex property that also depends on factors like lattice energy, entropy changes, and the nature of the solvent.

25. Why does atomic radius generally increase down a group in the periodic table?

Atomic radius generally increases down a group because each successive element has an additional electron shell. This new shell is further from the nucleus, increasing the overall size of the atom. The effect of increased nuclear charge is outweighed by the addition of a new shell.

26. How does atomic radius change across a period in the periodic table?

Atomic radius generally decreases across a period from left to right. This is because the number of protons (and electrons) increases, but electrons are being added to the same shell. The increased nuclear charge pulls the electrons closer to the nucleus, causing the atom to shrink.

27. What is the difference between atomic radius and ionic radius?

Atomic radius refers to the size of a neutral atom, while ionic radius is the size of an ion (a charged atom). Cations (positively charged ions) are typically smaller than their parent atoms because they've lost electrons. Anions (negatively charged ions) are usually larger than their parent atoms due to gained electrons.

28. What is the relationship between atomic radius and electronegativity?

There's generally an inverse relationship between atomic radius and electronegativity. As atomic radius decreases across a period, electronegativity increases. This is because smaller atoms with higher effective nuclear charge hold their electrons more tightly and are more likely to attract electrons from other atoms.

29. How does atomic radius affect chemical reactivity?

Atomic radius influences chemical reactivity by affecting an atom's ability to form bonds. Smaller atoms generally form stronger covalent bonds because their valence electrons are closer to the nucleus and held more tightly. Larger atoms tend to form weaker bonds but may be more reactive in ionic or metallic bonding scenarios.

30. What is the difference between covalent radius and van der Waals radius?

Covalent radius is half the distance between the nuclei of two covalently bonded atoms of the same element. Van der Waals radius is half the distance between the nuclei of two non-bonded atoms of the same element when they are as close as possible without forming a bond. Van der Waals radius is typically larger than covalent radius.

31. What is atomic radius?

Atomic radius is the distance from the center of an atom's nucleus to its outermost electron shell. It represents the size of an atom and is an important property that influences chemical behavior and bonding.

32. How does the concept of atomic radius relate to the octet rule?

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with eight valence electrons. This often involves changes in atomic size. For example, when sodium loses an electron to form a cation, its radius decreases significantly, while chlorine gains an electron to form a larger anion.

33. How do quantum mechanical models explain atomic radius trends?

Quantum mechanical models describe electrons as probability distributions (orbitals) rather than discrete particles. These models explain atomic radius trends through concepts like electron shielding, effective nuclear charge, and the aufbau principle. They account for the complex interplay between nuclear attraction and electron-electron repulsion that determines atomic size.

34. How does atomic radius affect the melting and boiling points of elements?

Generally, elements with larger atomic radii have lower melting and boiling points. This is because larger atoms typically have weaker interatomic forces, requiring less energy to overcome these forces and change state. However, other factors like electron configuration and types of bonding also play significant roles.

35. Why do elements in the same group have similar chemical properties despite having different atomic radii?

Elements in the same group have similar chemical properties primarily due to their similar valence electron configurations, not their atomic radii. While atomic size affects some properties, the number and arrangement of valence electrons largely determine chemical behavior. The similar outer electron structure leads to comparable reactivity patterns within a group.

36. Why do noble gases have relatively large atomic radii compared to other elements in the same period?

Noble gases have completely filled outer electron shells, which are more stable and experience less nuclear attraction than partially filled shells. This results in less electron compression and larger atomic radii compared to other elements in the same period with fewer electrons in their outer shell.

37. What is the lanthanide contraction, and how does it affect atomic radii?

The lanthanide contraction is the decrease in atomic and ionic radii across the lanthanide series. It occurs because the added electrons in the 4f subshell don't effectively shield the outer electrons from the increasing nuclear charge. This contraction affects the sizes of elements in subsequent periods, particularly the 6th and 7th periods.

38. Why do some elements, like mercury, have unusually small atomic radii?

Mercury's unusually small atomic radius is due to the lanthanide contraction and relativistic effects. The filled 4f subshell in lanthanides leads to increased nuclear charge without a corresponding increase in shielding. Additionally, the high speed of inner electrons in heavy atoms like mercury causes relativistic contraction of s orbitals, further reducing atomic size.

39. Why do transition metals often have similar atomic radii across a period?

Transition metals often have similar atomic radii across a period because they add electrons to inner d-orbitals rather than outer s-orbitals. The additional electrons and protons largely cancel each other out in terms of their effect on atomic size, resulting in only small changes in radius across the transition metal series.

40. How does atomic radius relate to ionization energy?

There's generally an inverse relationship between atomic radius and ionization energy. As atomic radius decreases, ionization energy typically increases. This is because smaller atoms hold their electrons more tightly due to greater effective nuclear charge, requiring more energy to remove an electron.

41. How does atomic radius affect the formation of metallic bonds?

Atomic radius influences metallic bonding by affecting how closely metal atoms can pack together. Smaller atoms generally form stronger metallic bonds because their valence electrons are closer to the nuclei, allowing for better overlap of electron clouds. This impacts properties like melting point, electrical conductivity, and malleability.

42. How does atomic radius relate to the concept of polarizability?

Polarizability generally increases with atomic radius. Larger atoms have more diffuse electron clouds that are farther from the nucleus and thus more easily distorted by external electric fields. This relationship between size and polarizability affects properties like van der Waals forces, refractive index, and some aspects of chemical reactivity.

43. Why do some pairs of elements, like nickel and palladium, have very similar atomic radii despite being in different periods?

Some elements in different periods can have similar atomic radii due to a combination of factors including the lanthanide contraction and the filling of inner d or f orbitals. In the case of nickel and palladium, the contraction of the 5s orbital in palladium (due to poor shielding by 4d electrons) roughly compensates for the additional electron shell, resulting in a similar radius to nickel.

44. What is the relationship between atomic radius and electron affinity?

There's generally an inverse relationship between atomic radius and electron affinity. As atomic radius decreases, electron affinity typically increases. This is because smaller atoms can more effectively attract and hold an additional electron due to the closer proximity of the nucleus and the increased effective nuclear charge.

45. How does atomic radius affect the formation of coordinate covalent bonds?

Atomic radius influences coordinate covalent (or dative) bonds by affecting the availability and accessibility of lone pairs of electrons. Smaller atoms with lone pairs (like nitrogen or oxygen) can form stronger coordinate bonds because their lone pairs are held more tightly and are in closer proximity to the accepting atom or ion.

46. Why do some elements, like the noble gases, not follow the general trend of decreasing atomic radius across a period?

Noble gases have fully filled outer electron shells, which are more stable and experience less nuclear attraction than partially filled shells. This results in less electron compression and larger atomic radii compared to the elements immediately preceding them in the period, creating a slight "bump" in the trend of decreasing radius across a period.

47. What is the significance of atomic radius in understanding the reactivity of halogens?

Atomic radius plays a crucial role in halogen reactivity. As you move down the group, the atomic radius increases, making it harder for the atom to attract electrons. This results in decreasing reactivity from fluorine to iodine. The smaller size of fluorine contributes to its extreme reactivity, as it can more easily attract electrons to complete its octet.

48. How does atomic radius affect the strength of intermolecular forces?

Atomic radius influences intermolecular forces, particularly dispersion forces (London forces). Larger atoms have more electrons and are more polarizable, leading to stronger dispersion forces. This is why, for instance, the boiling points of noble gases increase down the group as their atomic radii increase, despite their lack of permanent dipoles.

49. Why is there a larger jump in atomic radius between the 2nd and 3rd periods compared to later periods?

The larger jump in atomic radius between the 2nd and 3rd periods is due to the addition of a new principal quantum level (n=3). This represents a significant increase in the distance of electrons from the nucleus. In later periods, the effect of adding a new shell is partially offset by increased nuclear charge and the poor shielding of d and f electrons.

50. How does the concept of atomic radius relate to the formation of complex ions?

Atomic radius affects the formation of complex ions by influencing the ability of a central metal ion to accommodate ligands. Smaller metal ions can often form stronger bonds with ligands and accommodate fewer of them due to space constraints. Larger metal ions can often accommodate more ligands but may form weaker bonds. This impacts properties like stability constants and geometries of complex ions.

51. What is the relationship between atomic radius and the concept of electronegativity?

There's generally an inverse relationship between atomic radius and electronegativity. As atomic radius decreases, electronegativity typically increases. This is because smaller atoms have a greater effective nuclear charge, allowing them to more strongly attract electrons in a chemical bond. This relationship helps explain many trends in chemical bonding and reactivity.

52. How does atomic radius affect the stability of different oxidation states in transition metals?

Atomic radius influences the stability of oxidation states in transition metals. Generally, higher oxidation states are more stable for smaller atoms because they can better accommodate the increased positive charge. As atomic radius increases down a group