Misplaced Pages

#23976

59-441: 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

118-403: A nickel atom has, in principle, ten valence electrons (4s 3d ), its oxidation state never exceeds four. For zinc , the 3d subshell is complete in all known compounds, although it does contribute to the valence band in some compounds. Similar patterns hold for the ( n −2)f energy levels of inner transition metals. The d electron count is an alternative tool for understanding the chemistry of

177-477: 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. However, in group XIII ( Boron family ), the electronegativity first decreases from boron to aluminium and then increases down

236-448: 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. 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

295-406: A valence shell , thereby weakening 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 . The tendency of an atom in a molecule to attract the shared pair of electrons towards itself

354-439: A 1s configuration with two valence electrons, and thus having some similarities with the alkaline earth metals with their n s valence configurations, its shell is completely full and hence it is chemically very inert and is usually placed in group 18 with the other noble gases. The valence shell is the set of orbitals which are energetically accessible for accepting electrons to form chemical bonds . For main-group elements,

413-411: A closed shell are highly reactive due to the relatively low energy to remove the extra valence electrons to form a positive ion . An atom with one or two electrons fewer than a closed shell is reactive due to its tendency either to gain the missing valence electrons and form a negative ion, or else to share valence electrons and form a covalent bond. Similar to a core electron , a valence electron has

472-508: A core configuration identical to that of the noble gas argon . In this atom, a 3d electron has energy similar to that of a 4s electron, and much higher than that of a 3s or 3p electron. In effect, there are possibly seven valence electrons (4s 3d ) outside the argon-like core; this is consistent with the chemical fact that manganese can have an oxidation state as high as +7 (in the permanganate ion: MnO 4 ). (But note that merely having that number of valence electrons does not imply that

531-527: A covalent bond, one electron from the halogen and one electron from another atom form a shared pair (e.g., in the molecule H–F, the line represents a shared pair of valence electrons, one from H and one from F). Within each group of nonmetals, reactivity decreases with each lower row of the table (from a light element to a heavy element) in the periodic table, because the valence electrons are at progressively higher energies and thus progressively less tightly bound. In fact, oxygen (the lightest element in group 16)

590-440: A given element's reactivity is highly dependent upon its electronic configuration . For a main-group element , a valence electron can exist only in the outermost electron shell ; for a transition metal , a valence electron can also be in an inner shell. An atom with a closed shell of valence electrons (corresponding to a noble gas configuration ) tends to be chemically inert . Atoms with one or two valence electrons more than

649-402: A given element, but they are all at similar energies. As a general rule, a main-group element (except hydrogen or helium) tends to react to form a s p electron configuration . This tendency is called the octet rule , because each bonded atom has 8 valence electrons including shared electrons. Similarly, a transition metal tends to react to form a d s p electron configuration. This tendency

SECTION 10

#1732786929024

708-494: A heavier element), because a heavier element has more electron shells than a lighter element; a heavier element's valence electrons exist at higher principal quantum numbers (they are farther away from the nucleus of the atom, and are thus at higher potential energies, which means they are less tightly bound). A nonmetal atom tends to attract additional valence electrons to attain a full valence shell; this can be achieved in one of two ways: An atom can either share electrons with

767-476: A neighboring atom (a covalent bond ), or it can remove electrons from another atom (an ionic bond ). The most reactive kind of nonmetal element is a halogen (e.g., fluorine (F) or chlorine (Cl)). Such an atom has the following electron configuration: s p ; this requires only one additional valence electron to form a closed shell. To form an ionic bond, a halogen atom can remove an electron from another atom in order to form an anion (e.g., F , Cl , etc.). To form

826-434: A nonmetal; a semiconductor also differs from a metal in that a semiconductor's conductivity increases with temperature . The typical elemental semiconductors are silicon and germanium , each atom of which has four valence electrons. The properties of semiconductors are best explained using band theory , as a consequence of a small energy gap between a valence band (which contains the valence electrons at absolute zero) and

885-416: A significant increase in atomic radius with the first elements of each period. The atomic radius of each element generally decreases across each period due to an increasing number of protons, since an increase in the number of protons increases the attractive force acting on the atom's electrons. The greater attraction draws the electrons closer to the protons, decreasing the size of the atom. Down each group,

944-457: A sufficient fraction of the speed of light to gain a nontrivial amount of mass. The following table summarizes the main phenomena that influence the atomic radius of an element: 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

1003-412: A transition metal. The number of valence electrons of an element can be determined by the periodic table group (vertical column) in which the element is categorized. In groups 1–12, the group number matches the number of valence electrons; in groups 13–18, the units digit of the group number matches the number of valence electrons. (Helium is the sole exception.) Helium is an exception: despite having

1062-446: Is called the 18-electron rule , because each bonded atom has 18 valence electrons including shared electrons. The heavy group 2 elements calcium, strontium, and barium can use the ( n −1)d subshell as well, giving them some similarities to transition metals. The number of valence electrons in an atom governs its bonding behavior. Therefore, elements whose atoms have the same number of valence electrons are often grouped together in

1121-417: Is easily lost to form a positive ion (cation) with a closed shell (e.g., Na or K ). An alkaline earth metal of group 2 (e.g., magnesium ) is somewhat less reactive, because each atom must lose two valence electrons to form a positive ion with a closed shell (e.g., Mg ). Within each group (each periodic table column) of metals, reactivity increases with each lower row of the table (from a light element to

1180-490: Is found toward the right of the periodic table, and it has a valence shell that is at least half full (the exception is boron ). Its ionization energy is large; an electron cannot leave an atom easily when an electric field is applied, and thus such an element can conduct only very small electric currents. Examples of solid elemental insulators are diamond (an allotrope of carbon ) and sulfur . These form covalently bonded structures, either with covalent bonds extending across

1239-410: Is known as electron affinity. 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, as one moves down in a group , electron affinity decreases because atomic size increases due to the addition of

SECTION 20

#1732786929024

1298-482: Is known as electronegativity. It is a dimensionless quantity because it is only a tendency. 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 . Trend-wise, as one moves from left to right across

1357-437: Is known as the lanthanide contraction . A similar phenomenon exists for actinides ; however, the general instability of transuranic elements makes measurements for the remainder of the 5f-block difficult and for transactinides nearly impossible. Finally, for sufficiently heavy elements, the atomic radius may be decreased by relativistic effects . This is a consequence of electrons near the strongly charged nucleus traveling at

1416-484: 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

1475-420: Is not fully occupied. The electrons that determine valence – how an atom reacts chemically – are those with the highest energy . For a main-group element , the valence electrons are defined as those electrons residing in the electronic shell of highest principal quantum number n . Thus, the number of valence electrons that it may have depends on the electron configuration in a simple way. For example,

1534-402: Is noticeable up to platinum ( Z  = 78), after which it is masked by a relativistic effect known as the inert-pair effect . 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

1593-501: Is relatively free to leave one atom in order to associate with another nearby. This situation characterises metallic bonding . Such a "free" electron can be moved under the influence of an electric field , and its motion constitutes an electric current ; it is responsible for the electrical conductivity of the metal. Copper , aluminium , silver , and gold are examples of good conductors. A nonmetallic element has low electrical conductivity; it acts as an insulator . Such an element

1652-427: 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. 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

1711-417: Is the most reactive nonmetal after fluorine, even though it is not a halogen, because the valence shells of the heavier halogens are at higher principal quantum numbers. In these simple cases where the octet rule is obeyed, the valence of an atom equals the number of electrons gained, lost, or shared in order to form the stable octet. However, there are also many molecules that are exceptions , and for which

1770-467: 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. 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

1829-422: 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. 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

Periodic trends - Misplaced Pages Continue

1888-425: The n s level. So as opposed to main-group elements, a valence electron for a transition metal is defined as an electron that resides outside a noble-gas core. Thus, generally, the d electrons in transition metals behave as valence electrons although they are not in the outermost shell. For example, manganese (Mn) has configuration 1s 2s 2p 3s 3p 4s 3d ; this is abbreviated to [Ar] 4s 3d , where [Ar] denotes

1947-399: The periodic table of the elements, especially if they also have the same types of valence orbitals. The most reactive kind of metallic element is an alkali metal of group 1 (e.g., sodium or potassium ); this is because such an atom has only a single valence electron. During the formation of an ionic bond , which provides the necessary ionization energy , this one valence electron

2006-534: 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. 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,

2065-441: 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 . The concept of atomic radius

2124-406: The ability to absorb or release energy in the form of a photon . An energy gain can trigger the electron to move (jump) to an outer shell; this is known as atomic excitation . Or the electron can even break free from its associated atom's shell; this is ionization to form a positive ion. When an electron loses energy (thereby causing a photon to be emitted), then it can move to an inner shell which

2183-450: 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). The following table shows atomic radii computed from theoretical models, as published by Enrico Clementi and others in 1967. The values are in picometres (pm). Valence electron#Valence shell In chemistry and physics , valence electrons are electrons in

2242-430: The atomic radius of each element typically increases because there are more occupied electron energy levels and therefore a greater distance between protons and electrons. 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 despite an increase in attractive force from

2301-410: 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. The energy released when an electron is added to a neutral gaseous atom to form an anion

2360-495: The atomic size, getting 10 –10 cm for copper. The earliest estimates of the atomic size was made by opticians in the 1830s, particularly Cauchy , who developed models of light dispersion assuming a lattice of connected "molecules". In 1857 Clausius developed a gas-kinetic model which included the equation for mean free path . In the 1870s it was used to estimate gas molecule sizes, as well as an aforementioned comparison with visible light wavelength and an estimate from

2419-406: 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 ), and less than 1/1000 of

Periodic trends - Misplaced Pages Continue

2478-456: The box ranges from red to yellow as the radius increases; gray indicates lack of data. Electrons in atoms fill electron shells from the lowest available energy level. As a consequence of the Aufbau principle , each new period begins with the first two elements filling the next unoccupied s-orbital . Because an atom's s-orbital electrons are typically farthest from the nucleus, this results in

2537-436: The corresponding oxidation state will exist. For example, fluorine is not known in oxidation state +7; and although the maximum known number of valence electrons is 16 in ytterbium and nobelium , no oxidation state higher than +9 is known for any element.) The farther right in each transition metal series, the lower the energy of an electron in a d subshell and the less such an electron has valence properties. Thus, although

2596-458: The electronic configuration of phosphorus (P) is 1s 2s 2p 3s 3p so that there are 5 valence electrons (3s 3p ), corresponding to a maximum valence for P of 5 as in the molecule PF 5 ; this configuration is normally abbreviated to [Ne] 3s 3p , where [Ne] signifies the core electrons whose configuration is identical to that of the noble gas neon . However, transition elements have ( n −1)d energy levels that are very close in energy to

2655-468: 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 . The valency of an element is the number of electrons that must be lost or gained by an atom to obtain

2714-403: 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. Atomic radius 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

2773-413: The most stable allotrope is considered. Metallic elements generally have high electrical conductivity when in the solid state. In each row of the periodic table , the metals occur to the left of the nonmetals, and thus a metal has fewer possible valence electrons than a nonmetal. However, a valence electron of a metal atom has a small ionization energy , and in the solid-state this valence electron

2832-407: The nucleus. Electron shielding causes the attraction of an atom's nucleus on its electrons to decrease, so electrons occupying higher energy states farther from the nucleus experience reduced attractive force, increasing the size of the atom. However, elements in the 5d-block ( lutetium to mercury ) are much smaller than this trend predicts due to the weak shielding of the 4f-subshell. This phenomenon

2891-462: 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 presence of valence electrons can determine the element 's chemical properties, such as its valence —whether it may bond with other elements and, if so, how readily and with how many. In this way,

2950-402: 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

3009-412: The radius may depend on the atom's state and context. 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

SECTION 50

#1732786929024

3068-464: The thickness of soap bubble film at which its contractile force rapidly diminishes. By 1900, various estimates of mercury atom diameter averaged around 275±20 pm (modern estimates give 300±10 pm, see below). 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. However, in 1923, when more crystal data had become available, it

3127-410: The valence electrons of the metal atoms are used to form ionic bonds . For example, although elemental sodium is a metal, solid sodium chloride is an insulator, because the valence electron of sodium is transferred to chlorine to form an ionic bond, and thus that electron cannot be moved easily. A semiconductor has an electrical conductivity that is intermediate between that of a metal and that of

3186-409: The valence is less clearly defined. Valence electrons are also responsible for the bonding in the pure chemical elements, and whether their electrical conductivity is characteristic of metals, semiconductors, or insulators. Metallic Network covalent Molecular covalent Single atoms Unknown Background color shows bonding of simple substances in the periodic table . If there are several,

3245-404: The valence shell consists of the n s and n p orbitals in the outermost electron shell . For transition metals the orbitals of the incomplete ( n −1)d subshell are included, and for lanthanides and actinides incomplete ( n −2)f and ( n −1)d subshells. The orbitals involved can be in an inner electron shell and do not all correspond to the same electron shell or principal quantum number n in

3304-400: 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

3363-433: The whole structure (as in diamond) or with individual covalent molecules weakly attracted to each other by intermolecular forces (as in sulfur). (The noble gases remain as single atoms, but those also experience intermolecular forces of attraction, that become stronger as the group is descended: helium boils at −269 °C, while radon boils at −61.7 °C.) A solid compound containing metals can also be an insulator if

3422-419: Was found that the approximation of an atom as a sphere does not necessarily hold when comparing the same atom in different crystal structures. 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. The values are in picometers (pm or 1×10  m), with an accuracy of about 5 pm. The shade of

3481-450: Was preceded in the 19th century by the concept of atomic volume, a relative measure of how much space would on average an atom occupy in a given solid or liquid material. By the end of the century this term was also used in an absolute sense, as a molar volume divided by Avogadro constant . Such a volume is different for different crystalline forms even of the same compound, but physicists used it for rough, order-of-magnitude estimates of

#23976