Name: periodic trends in atomic size
Uploaded: 2020-09-03T13:38:28.000Z
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Atomic Radii and Effective Nuclear Charge

La taille d'un atome est dictée par les électrons ou leurs orbitales. Cependant, les orbitales ne décrivent pas un espace confiné, mais plutôt la probabilité statistique de la localisation d'un électron Alors, comment la taille atomique est-elle définie et qu'est-ce qui l'influence? Un rayon atomique peut être décrit de deux manières.Le rayon atomique non lié, ou rayon de van der Waals d'un atome, est la moitié de la distance entre les noyaux adjacents dans le solide atomique. Inversement, un rayon atomique de liaison, ou rayon covalent, distingue les métaux et les non-métaux. Dans les métaux, le rayon est décrit pour les atomes dans leur structure cristalline comme la moitié de la distance entre les centres de deux atomes voisins.Dans les molécules diatomiques non métalliques, le rayon est décrit comme la moitié de la distance entre les centres des atomes liés. Le tableau périodique représente les variations du rayonnement covalent qui sont appelés rayons atomiques, qui sont influencés par deux facteurs;le nombre de niveaux d'énergie principaux des électrons de valence, et la charge nucléaire effective. La tendance des rayons atomiques pour les éléments du groupe principal en bas des colonnes qui est représentée ici.En descendant un groupe, le nombre quantique principal, n, augmente de un pour chaque élément. Ainsi, alors que les électrons externes s'éloignent du noyau, le rayon atomique augmente dans le groupe. Par exemple, en descendant du groupe 1, le rayon atomique augmente du lithium au césium.Cette tendance est démontrée par l'ensemble du tableau périodique. De plus, l'intrigue révèle que le rayon atomique est au maximum pour chaque métal alcalin et tombe, au minimum par chaque gaz noble au cours de la période. La diminution des rayons atomiques sur une période peut s'expliquer par la charge nucléaire effective.Rappelons le concept d'une charge nucléaire efficace. Dans tout atome à plusieurs électrons, les électrons de la couche intrne protègent partiellement les électrons de la couche externe de la traction du noyau. Ainsi, la charge nucléaire effective, la charge ressentie par un électron externe est inférieure à la charge nucléaire réelle.Les électrons dans la même couche de valence ne se protègent pas très efficacement. Au cours de la période, la charge nucléaire augmente tandis que le nombre d'électrons de la couche interne reste constant. Ainsi, à mesure que la charge nucléaire effective augmente régulièrement, le blindage des électrons externes diminue, ce qui entraîne une diminution des rayons atomiques.Les rayons de la plupart des éléments de transition, cependant, restent à peu près constants à travers chaque rangée. En effet, le nombre d'électrons dans le niveau d'énergie principal externe est presque constant.

rayons atomiques et charge nucléaire effective

Les éléments des groupes du tableau périodique présentent un comportement chimique similaire entre eux. Cette similarité existe parce que les membres d'un groupe ont le même nombre et la même distribution d'électrons sur leurs couches de valence. 
 
 Sur une période de gauche à droite, un proton est ajouté au noyau et un électron est ajouté à la couche de valence avec chaque élément successif. En descendant le long des éléments d'un groupe, le nombre d'électrons sur la couche de valence reste constant, mais le nombre quantique principal augmente de un à chaque fois. Comprendre la structure électronique des éléments nous permet d'examiner certaines des propriétés qui régissent leur comportement chimique. Ces propriétés varient de façon périodique à mesure que la structure électronique des éléments change.  
 
 <h4>Variation du rayon atomique </h4>
 
 À cause de l'image de la mécanique quantique, il est difficile d'établir la taille précise d'un atome. Cependant, il existe plusieurs moyens pratiques de définir le rayon des atomes et, par conséquent, de déterminer leurs tailles relatives qui donnent des valeurs à peu près similaires. 
 
 Le rayon atomique des métaux représente la moitié de la distance entre les centres de deux atomes voisins. C'est la moitié de la distance entre les centres des atomes liés pour les éléments qui existent en tant que molécules diatomiques. 
 
 En se déplaçant de gauche à droite sur une période, généralement, chaque élément a un rayon atomique plus petit que l'élément qui le précède. Cela peut sembler contre-intuitif car cela implique que les atomes ayant plus d'électrons ont un rayon atomique plus petit. On peut expliquer cela en se basant sur l'idée d'une charge nucléaire effective. Dans n'importe quel atome à plusieurs électrons, les électrons sur la couche interne protègent partiellement les électrons sur la couche externe de l'attraction du noyau.  Ainsi, la charge nucléaire effective, la charge ressentie par un électron, est inférieure à la charge nucléaire réelle (Z) et peut être estimée comme suit : 
 
 Zeff = Z – σ                
 
 où, Zeff est la charge nucléaire effective, Z la charge nucléaire réelle et σ la constante de l'effet d'écran, où la constante de l'effet d'écran est supérieure à zéro mais inférieure à Z. 
 
 Chaque fois que nous passons d'un élément à l'autre le long d'une période, Z augmente de un, mais l'effet d'écran n'augmente que légèrement. Ainsi, Zeff augmente à mesure que nous nous déplaçons de gauche à droite sur une période. L'attraction plus forte (charge nucléaire effective plus élevée) que connaissent les électrons sur le côté droit du tableau périodique les rapproche du noyau, ce qui réduit les rayons atomiques. 
 
 Les électrons de cœur protègent efficacement les électrons du niveau principal le plus à l'extérieur de la charge nucléaire, mais les électrons les plus à l'extérieur ne se protègent pas les uns les autres de la charge nucléaire de manière efficace. Plus la charge nucléaire effective est grande, plus l'emprise du noyau sur les électrons externes est forte et plus le rayon atomique est petit. 
 
 Cependant, les rayons de certains éléments de transition restent à peu près constants sur chaque ligne. En effet, le nombre d'électrons sur le niveau d'énergie principal le plus externe est presque constant et ils subissent une charge nucléaire effective à peu près constante. 
 
 Dans chaque période, la tendance est que le rayon atomique diminue à mesure que Z augmente ;  dans chaque groupe, la tendance est que le rayon atomique augmente à mesure que Z augmente. 
 
 En descendant dans un groupe, le nombre quantique principal, n, augmente de un pour chaque élément. Ainsi, les électrons sont ajoutés à une région de l'espace de plus en plus éloignée du noyau. Par conséquent, la taille de l'atome (et son rayon atomique) doit augmenter à mesure que nous augmentons la distance entre les électrons les plus à l'extérieur du noyau. Cette tendance est illustrée pour les rayons atomiques des halogènes dans le tableau ci-dessous. 
 
 <table style="margin-left: auto; margin-right: auto;">
 <tbody style="border: 1px solid black;">
 <tr>
 <td colspan="3" style="color: black; border: 1px solid black; text-align: center;">Rayons atomiques des éléments du groupe des halogènes </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Atome </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Rayon atomique (pm) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Charge nucléaire, Z</td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">F </td>
 <td style="color: black; border: 1px solid black; text-align: center;">64 </td>
 <td style="color: black; border: 1px solid black; text-align: center;">9+ </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Cl </td>
 <td style="color: black; border: 1px solid black; text-align: center;">99 </td>
 <td style="color: black; border: 1px solid black; text-align: center;">17+ </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Br </td>
 <td style="color: black; border: 1px solid black; text-align: center;">114 </td>
 <td style="color: black; border: 1px solid black; text-align: center;">35+ </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">I </td>
 <td style="color: black; border: 1px solid black; text-align: center;">133 </td>
 <td style="color: black; border: 1px solid black; text-align: center;">53+ </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">At </td>
 <td style="color: black; border: 1px solid black; text-align: center;">148 </td>
 <td style="color: black; border: 1px solid black; text-align: center;">85+ </td>
 </tr>
 </tbody>
 </table>
 
  
 
 Ce texte est adapté de <a href="https://openstax.org/books/chemistry-2e/pages/6-5-periodic-variations-in-element-properties">Openstax Chemistry 2e, Section 6.5 : Variations périodiques dans les propriétés des éléments.</a>

Rayons atomiques et charge nucléaire effective

Les éléments des groupes du tableau périodique présentent un comportement chimique similaire. Cette similitude se produit parce que les membres d’un groupe ont le même nombre ...

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periodic trends in atomic size

Periodic trends in atomic size help explain why atoms get larger or smaller across the periodic table. Elements in the same group show similar chemical behavior because they have the same number and arrangement of valence electrons. Across a period, each new element adds one proton to the nucleus and one electron to the valence shell. Down a group, the number of valence electrons stays the same, but the principal quantum number increases by one.

Atomic radius is not easy to define with perfect precision because atoms do not have a hard edge. Scientists use practical measurements instead. In metals, atomic radius is one-half the distance between the centers of two neighboring atoms. For elements that form diatomic molecules, it is one-half the distance between the centers of bonded atoms.

Moving from left to right across a period, atomic radius generally decreases. That may seem surprising because atoms on the right have more electrons. The reason is effective nuclear charge, or the charge an electron feels from the nucleus. It can be estimated as Z eff = Z – σ, where Z is the actual nuclear charge and σ is the shielding constant.

As Z increases by one from element to element across a period, shielding rises only a little. This means Z eff also increases from left to right. A stronger pull from the nucleus draws outer electrons closer, so the atomic radius becomes smaller. Core electrons shield outer electrons well, but outermost electrons do not shield one another very effectively. Some transition elements keep nearly the same radius across a row because the number of electrons in the outermost principal energy level changes very little, so the effective nuclear charge stays roughly constant.

The trend within a group is the opposite. Atomic radius increases as Z increases down a group. Each step down adds a principal energy level, so the outermost electrons are farther from the nucleus. That greater distance makes the atom larger. The halogens show this pattern clearly, with atomic radii increasing from fluorine to astatine.

Periodic Trends in Atomic Size

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Ionization energy is the energy needed to remove an electron from a gaseous atom or ion in its ground state. The first ionization energy, or IE1, removes the most loosely bound electron and forms a 1+ cation. The second ionization energy, IE2, removes the next electron, and each later ionization energy removes another electron from the atom or ion. Ionization always requires energy, so these processes are endothermic. That means ionization energy values are always positive. In general, larger...

atomic radii and effective nuclear charge

Periodic trends in atomic size help explain why atoms get larger or smaller across the periodic table. Elements in the same group show similar chemical behavior because they have the same number and arrangement of valence electrons. Across a period, each new element adds one proton to the nucleus and one electron to the valence shell. Down a group, the number of valence electrons stays the same, but the principal quantum number increases by one. Atomic radius is not easy to define with perfect...

ionic radii

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Ionic radius describes the size of an ion. It changes when an atom gains or loses electrons. Cations are positive ions with fewer electrons than the parent atom, and they are smaller than the atoms they come from. Aluminum shows this pattern clearly. A neutral aluminum atom has a covalent radius of 118 pm, while Al3+ has an ionic radius of 68 pm. When electrons are removed from the outer valence shell, the remaining core electrons feel a stronger effective nuclear charge, or Zeff. That pull...

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Alkali metals are Group 1 elements with one valence electron. They are soft, shiny metallic solids that are malleable, ductile, and good conductors of heat and electricity. Their melting points are unusually low for metals and drop as you move down the group. Density generally increases down the group, except for potassium. The trend in atomic size helps explain their behavior. Moving from lithium to cesium, each element has a higher principal quantum number, so the outer electron is farther...

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Halogens are Group 17 nonmetals that show clear trends in structure and reactivity. At room temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. Astatine is highly unstable and radioactive, so most of its properties are still unknown because it has a very short half-life. Tennessine is a synthetic element that is also predicted to belong to this group. These elements do not usually exist as single atoms. Instead, they form diatomic molecules, which means...

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Noble gases are the group 18 elements helium, neon, argon, krypton, xenon, and radon. They were called “noble” because they were once thought to be nonreactive. That idea changed in 1962, when Dr. Neil Bartlett at the University of British Columbia showed that this assumption was not fully correct. These gases occur in very small amounts in the atmosphere. Some natural gas contains 1–2% helium by mass. Helium is separated from natural gas by removing the condensable components, which leaves...

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Learn why atomic radius changes across periods and groups, and how effective nuclear charge and shielding shape the size of atoms in the periodic table.

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Periodic Trends in Atomic Size

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