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.
Depending on the definition, the term may apply to atoms in condensed matter, covalently bonding in molecules, or in ionized and excited states; and its value may be obtained through experimental measurements, or computed from theoretical models. The value of the radius may depend on the atom's state and context. [1]
Electrons do not have definite orbits nor sharply defined ranges. Rather, their positions must be described as probability distributions that taper off gradually as one moves away from the nucleus, without a sharp cutoff; these are referred to as atomic orbitals or electron clouds. Moreover, in condensed matter and molecules, the electron clouds of the atoms usually overlap to some extent, and some of the electrons may roam over a large region encompassing two or more atoms.
Under most definitions the radii of isolated neutral atoms range between 30 and 300 pm (trillionths of a meter), or between 0.3 and 3 ångströms. Therefore, the radius of an atom is more than 10,000 times the radius of its nucleus (1–10 fm), [2] and less than 1/1000 of the wavelength of visible light (400–700 nm).
For many purposes, atoms can be modeled as spheres. This is only a crude approximation, but it can provide quantitative explanations and predictions for many phenomena, such as the density of liquids and solids, the diffusion of fluids through molecular sieves, the arrangement of atoms and ions in crystals, and the size and shape of molecules.[ citation needed ]
In 1920, shortly after it had become possible to determine the sizes of atoms using X-ray crystallography, it was suggested that all atoms of the same element have the same radii. [3] However, in 1923, when more crystal data had become available, it was found that the approximation of an atom as a sphere does not necessarily hold when comparing the same atom in different crystal structures. [4]
Widely used definitions of atomic radius include:
The following table shows empirically measured covalent radii for the elements, as published by J. C. Slater in 1964. [9] The values are in picometers (pm or 1×10−12 m), with an accuracy of about 5 pm. The shade of the box ranges from red to yellow as the radius increases; gray indicates lack of data.
Group (column) | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 | |||
Period (row) | |||||||||||||||||||||
1 | H 25 | He | |||||||||||||||||||
2 | Li 145 | Be 105 | B 85 | C 70 | N 65 | O 60 | F 50 | Ne | |||||||||||||
3 | Na 180 | Mg 150 | Al 125 | Si 110 | P 100 | S 100 | Cl 100 | Ar | |||||||||||||
4 | K 220 | Ca 180 | Sc 160 | Ti 140 | V 135 | Cr 140 | Mn 140 | Fe 140 | Co 135 | Ni 135 | Cu 135 | Zn 135 | Ga 130 | Ge 125 | As 115 | Se 115 | Br 115 | Kr | |||
5 | Rb 235 | Sr 200 | Y 180 | Zr 155 | Nb 145 | Mo 145 | Tc 135 | Ru 130 | Rh 135 | Pd 140 | Ag 160 | Cd 155 | In 155 | Sn 145 | Sb 145 | Te 140 | I 140 | Xe | |||
6 | Cs 260 | Ba 215 | * | Lu 175 | Hf 155 | Ta 145 | W 135 | Re 135 | Os 130 | Ir 135 | Pt 135 | Au 135 | Hg 150 | Tl 190 | Pb 180 | Bi 160 | Po 190 | At | Rn | ||
7 | Fr | Ra 215 | ** | Lr | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | Cn | Nh | Fl | Mc | Lv | Ts | Og | ||
* | La 195 | Ce 185 | Pr 185 | Nd 185 | Pm 185 | Sm 185 | Eu 185 | Gd 180 | Tb 175 | Dy 175 | Ho 175 | Er 175 | Tm 175 | Yb 175 | |||||||
** | Ac 195 | Th 180 | Pa 180 | U 175 | Np 175 | Pu 175 | Am 175 | Cm | Bk | Cf | Es | Fm | Md | No | |||||||
The way the atomic radius varies with increasing atomic number can be explained by the arrangement of electrons in shells of fixed capacity. The shells are generally filled in order of increasing radius, since the negatively charged electrons are attracted by the positively charged protons in the nucleus. As the atomic number increases along each row of the periodic table, the additional electrons go into the same outermost shell; whose radius gradually contracts, due to the increasing nuclear charge. In a noble gas, the outermost shell is completely filled; therefore, the additional electron of next alkali metal will go into the next outer shell, accounting for the sudden increase in the atomic radius.
The increasing nuclear charge is partly counterbalanced by the increasing number of electrons, a phenomenon that is known as shielding; which explains why the size of atoms usually increases down each column. However, there is one notable exception, known as the lanthanide contraction: the 5d block of elements are much smaller than one would expect, due to the weak shielding of the 4f electrons.
Essentially, the atomic radius decreases across the periods due to an increasing number of protons. Therefore, there is a greater attraction between the protons and electrons because opposite charges attract, and more protons create a stronger charge. The greater attraction draws the electrons closer to the protons, decreasing the size of the particle. Therefore, the atomic radius decreases. Down the groups, atomic radius increases. This is because there are more energy levels and therefore a greater distance between protons and electrons. In addition, electron shielding causes attraction to decrease, so remaining electrons can go farther away from the positively charged nucleus. Therefore, the size, or atomic radius, increases.
The following table summarizes the main phenomena that influence the atomic radius of an element:
factor | principle | increase with... | tend to | effect on radius |
---|---|---|---|---|
electron shells | quantum mechanics | principal and azimuthal quantum numbers | increase down each column | increases the atomic radius |
nuclear charge | attractive force acting on electrons by protons in nucleus | atomic number | increase along each period (left to right) | decreases the atomic radius |
shielding | repulsive force acting on outermost shell electrons by inner electrons | number of electrons in inner shells | reduce the effect of nuclear charge | increases the atomic radius |
The electrons in the 4f-subshell, which is progressively filled from lanthanum ( Z = 57) to ytterbium (Z = 70), are not particularly effective at shielding the increasing nuclear charge from the sub-shells further out. The elements immediately following the lanthanides have atomic radii which are smaller than would be expected and which are almost identical to the atomic radii of the elements immediately above them. [10] Hence lutetium is in fact slightly smaller than yttrium, hafnium has virtually the same atomic radius (and chemistry) as zirconium, and tantalum has an atomic radius similar to niobium, and so forth. The effect of the lanthanide contraction is noticeable up to platinum (Z = 78), after which it is masked by a relativistic effect known as the inert-pair effect.[ citation needed ]
Due to lanthanide contraction, the 5 following observations can be drawn:
The d-block contraction is less pronounced than the lanthanide contraction but arises from a similar cause. In this case, it is the poor shielding capacity of the 3d-electrons which affects the atomic radii and chemistries of the elements immediately following the first row of the transition metals, from gallium (Z = 31) to bromine (Z = 35). [10]
The following table shows atomic radii computed from theoretical models, as published by Enrico Clementi and others in 1967. [11] The values are in picometres (pm).
Group (column) | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 | ||
Period (row) | ||||||||||||||||||||
1 | H 53 | He 31 | ||||||||||||||||||
2 | Li 167 | Be 112 | B 87 | C 67 | N 56 | O 48 | F 42 | Ne 38 | ||||||||||||
3 | Na 190 | Mg 145 | Al 118 | Si 111 | P 98 | S 88 | Cl 79 | Ar 71 | ||||||||||||
4 | K 243 | Ca 194 | Sc 184 | Ti 176 | V 171 | Cr 166 | Mn 161 | Fe 156 | Co 152 | Ni 149 | Cu 145 | Zn 142 | Ga 136 | Ge 125 | As 114 | Se 103 | Br 94 | Kr 88 | ||
5 | Rb 265 | Sr 219 | Y 212 | Zr 206 | Nb 198 | Mo 190 | Tc 183 | Ru 178 | Rh 173 | Pd 169 | Ag 165 | Cd 161 | In 156 | Sn 145 | Sb 133 | Te 123 | I 115 | Xe 108 | ||
6 | Cs 298 | Ba 253 | * | Lu 217 | Hf 208 | Ta 200 | W 193 | Re 188 | Os 185 | Ir 180 | Pt 177 | Au 174 | Hg 171 | Tl 156 | Pb 154 | Bi 143 | Po 135 | At 127 | Rn 120 | |
7 | Fr | Ra | ** | Lr | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | Cn | Nh | Fl | Mc | Lv | Ts | Og | |
* | La 226 | Ce 210 | Pr 247 | Nd 206 | Pm 205 | Sm 238 | Eu 231 | Gd 233 | Tb 225 | Dy 228 | Ho 226 | Er 226 | Tm 222 | Yb 222 | ||||||
** | Ac | Th | Pa | U | Np | Pu | Am | Cm | Bk | Cf | Es | Fm | Md | No |
The atomic number or nuclear charge number (symbol Z) of a chemical element is the charge number of an atomic nucleus. For ordinary nuclei composed of protons and neutrons, this is equal to the proton number (np) or the number of protons found in the nucleus of every atom of that element. The atomic number can be used to uniquely identify ordinary chemical elements. In an ordinary uncharged atom, the atomic number is also equal to the number of electrons.
Atoms are the basic particles of the chemical elements. An atom consists of a nucleus of protons and generally neutrons, surrounded by an electromagnetically bound swarm of electrons. The chemical elements are distinguished from each other by the number of protons that are in their atoms. For example, any atom that contains 11 protons is sodium, and any atom that contains 29 protons is copper. Atoms with the same number of protons but a different number of neutrons are called isotopes of the same element.
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.
Ionic bonding is a type of chemical bonding that involves the electrostatic attraction between oppositely charged ions, or between two atoms with sharply different electronegativities, and is the primary interaction occurring in ionic compounds. It is one of the main types of bonding, along with covalent bonding and metallic bonding. Ions are atoms with an electrostatic charge. Atoms that gain electrons make negatively charged ions. Atoms that lose electrons make positively charged ions. This transfer of electrons is known as electrovalence in contrast to covalence. In the simplest case, the cation is a metal atom and the anion is a nonmetal atom, but these ions can be more complex, e.g. molecular ions like NH+
4 or SO2−
4. In simpler words, an ionic bond results from the transfer of electrons from a metal to a non-metal to obtain a full valence shell for both atoms.
A molecule is a group of two or more atoms held together by attractive forces known as chemical bonds; depending on context, the term may or may not include ions which satisfy this criterion. In quantum physics, organic chemistry, and biochemistry, the distinction from ions is dropped and molecule is often used when referring to polyatomic ions.
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.
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
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.
An extended periodic table theorizes about chemical elements beyond those currently known and proven. The element with the highest atomic number known is oganesson (Z = 118), which completes the seventh period (row) in the periodic table. All elements in the eighth period and beyond thus remain purely hypothetical.
Relativistic quantum chemistry combines relativistic mechanics with quantum chemistry to calculate elemental properties and structure, especially for the heavier elements of the periodic table. A prominent example is an explanation for the color of gold: due to relativistic effects, it is not silvery like most other metals.
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.
Ionic radius, rion, is the radius of a monatomic ion in an ionic crystal structure. Although neither atoms nor ions have sharp boundaries, they are treated as if they were hard spheres with radii such that the sum of ionic radii of the cation and anion gives the distance between the ions in a crystal lattice. Ionic radii are typically given in units of either picometers (pm) or angstroms (Å), with 1 Å = 100 pm. Typical values range from 31 pm (0.3 Å) to over 200 pm (2 Å).
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.
In molecular geometry, bond length or bond distance is defined as the average distance between nuclei of two bonded atoms in a molecule. It is a transferable property of a bond between atoms of fixed types, relatively independent of the rest of the molecule.
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 covalent radius of fluorine is a measure of the size of a fluorine atom; it is approximated at about 60 picometres.
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.
This glossary of chemistry terms is a list of terms and definitions relevant to chemistry, including chemical laws, diagrams and formulae, laboratory tools, glassware, and equipment. Chemistry is a physical science concerned with the composition, structure, and properties of matter, as well as the changes it undergoes during chemical reactions; it features an extensive vocabulary and a significant amount of jargon.
The atomic nucleus is the small, dense region consisting of protons and neutrons at the center of an atom, discovered in 1911 by Ernest Rutherford based on the 1909 Geiger–Marsden gold foil experiment. After the discovery of the neutron in 1932, models for a nucleus composed of protons and neutrons were quickly developed by Dmitri Ivanenko and Werner Heisenberg. An atom is composed of a positively charged nucleus, with a cloud of negatively charged electrons surrounding it, bound together by electrostatic force. Almost all of the mass of an atom is located in the nucleus, with a very small contribution from the electron cloud. Protons and neutrons are bound together to form a nucleus by the nuclear force.