Name: periodic trends in atomic size
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 3 min 8 s
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North Carolina K-12 Science Standard Course of Study

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Atomic Radii and Effective Nuclear Charge

O tamanho de um átomo é ditado pelos elétrons ou pelas suas orbitais. Contudo, as orbitais não descrevem um espaço confinado, mas sim a probabilidade estatística de onde se pode encontrar um elétron. Assim, como é definido o tamanho atômico e o que o influencia, um raio atômico pode ser descrito de duas maneiras.O raio atômico sem ligação ou o raio van der Waals de um átomo, é a metade da distância entre os núcleos adjacentes no sólido atômico. Inversamente, um raio atômico de ligação ou raio covalente, faz a distinção entre os metais e não metais. Nos metais, o raio é descrito para os átomos na sua estrutura de cristal como metade da distância entre os centros de dois átomos vizinhos.Nos não-metálicos, moléculas diatômicas, o raio é descrito como metade da distância entre os centros de átomos ligados. A tabela periódica apresenta variações nos raios covalentes que são frequentemente chamados de raios atômicos e são influenciados por dois fatores;o número dos principais níveis de energia de energia dos elétrons de valência e a carga nuclear efetiva. A tendência nos raios atômicos para os elementos do grupo principal das colunas está aqui representada.No grupo abaixo, o número quântico principal, n, aumenta em um para cada elemento. Assim, à medida que os elétrons exteriores se afastam mais do núcleo, o raio atômico aumenta pelo grupo abaixo. Por exemplo, descendo para o grupo 1, o raio atômico aumenta do lítio até ao césio.Esta tendência é demonstrada em toda a tabela periódica. Para além disso, a parcela revela que o raio atômico é o máximo para cada metal alcalino e cai a um mínimo com cada gás nobre em todo período. Os raios atômicos decrescentes ao longo de um período podem ser explicados pela carga nuclear efetiva.Recordemos o conceito de uma carga nuclear eficaz. Em qualquer átomo com múltiplos elétrons, os elétrons da camada interior da camada interior protegem parcialmente os elétrons da camada exterior a partir do centro do núcleo. Assim, a carga nuclear efetiva, a carga sentida por um elétron exterior é menor do que a carga nuclear real.Os elétrons na mesma camada de valência não se protegem uns aos outros de forma muito eficaz. Ao longo do período, a carga nuclear aumenta enquanto o número de elétrons da camada interna permanece constante. Assim, à medida que a carga nuclear efetiva aumenta de forma constante, a proteção dos elétrons exteriores torna-se menor, e isto leva a uma diminuição dos raios atômicos.Os raios da maioria dos elementos de transição, no entanto, mantêm-se aproximadamente constantes em cada fila. Isto porque o número de elétrons no nível energético principal mais periférico é quase constante.

radiação atómica e carga nuclear efetiva

Os elementos em grupos da tabela periódica apresentam comportamento químico semelhante. Essa semelhança ocorre porque os membros de um grupo têm o mesmo número e distribuição de eletrões nas suas camadas de valência.

Passando por um período da esquerda para a direita, um protão é adicionado ao núcleo e um eletrão à camada de valência a cada elemento sucessivo. Descendo os elementos de um grupo, o número de eletrões na camada de valência permanece constante, mas o número quântico principal aumenta em um de cada vez. Uma compreensão da estrutura eletrónica dos elementos permite-nos examinar algumas das propriedades que governam o seu comportamento químico. Estas propriedades variam periodicamente à medida que a estrutura eletrónica dos elementos muda. 

<h4>Variação no Raio Atómico</h4>

O quadro mecânico-quântico torna difícil estabelecer um tamanho definido de um átomo. No entanto, existem várias formas práticas de definir o raio dos átomos e, assim, determinar os seus tamanhos relativos que dão valores aproximadamente semelhantes.

O raio atómico em metais é metade da distância entre os centros de dois átomos vizinhos. É metade da distância entre os centros de átomos ligados para elementos que existem como moléculas diatómicas.

Movendo em um período da esquerda para a direita, geralmente, cada elemento tem um raio atómico menor do que o elemento que o precede. Isto pode parecer contra-intuitivo porque implica que átomos com mais eletrões têm um raio atómico menor. Isto pode ser explicado com base no conceito de carga nuclear efetiva. Em qualquer átomo com múltiplos eletrões, a camada interior de eletrões protege parcialmente camada exterior de eletrões da atração do núcleo.  Assim, a carga nuclear efetiva, a carga sentida por um eletrão, é menor do que a carga nuclear real (Z) e pode ser estimada pelo seguinte:

Zeff = Z – σ              

onde, Zeff é a carga nuclear efetiva, Z é a carga nuclear real, e σ é a constante de proteção, em que a constante de proteção é maior que zero, mas menor que Z.

Cada vez que nos movemos de um elemento para o outro ao longo de um período, Z aumenta em um, mas a a proteção aumenta apenas ligeiramente. Assim, Zeff aumenta à medida que nos movemos da esquerda para a direita ao longo de um período. A força de atração (carga nuclear efetiva mais alta) experienciada por eletrões do lado direito da tabela periódica aproxima-os do núcleo, tornando os raios atómicos mais pequenos.

Os eletrões centrais protegem eficazmente os eletrões no nível principal mais externo da carga nuclear, mas os eletrões mais externos não se protegem eficazmente uns dos outros da carga nuclear. Quanto maior a carga nuclear efetiva, mais fortemente o núcleo retém os eletrões externos, e menor o raio atómico.

No entanto, os raios de alguns elementos de transição permanecem aproximadamente constantes em cada linha. Isto deve-se ao facto de o número de eletrões no nível energético principal mais externo ser quase constante e de estes últimos terem uma carga nuclear efetiva aproximadamente constante.

Dentro de cada período, a tendência no raio atómico diminui à medida que Z aumenta; Dentro de cada grupo, a tendência é que o raio atómico aumente à medida que Z aumenta.

Descendo em um grupo, o número quântico principal, n, aumenta em um por cada elemento. Portanto, os eletrões estão a ser adicionados a uma região de espaço cada vez mais distante do núcleo. Consequentemente, o tamanho do átomo (e o seu raio atómico) deve aumentar à medida que aumentamos a distância dos eletrões mais externos do núcleo. Esta tendência é ilustrada para os raios atómicos dos halogénios na tabela abaixo.

<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;">Raios Atómicos dos Elementos do Grupo dos Halogéneos</td>
</tr>
<tr>
<td style="color: black; border: 1px solid black; text-align: center;">Átomo</td>
<td style="color: black; border: 1px solid black; text-align: center;">Raio atómico (pm)</td>
<td style="color: black; border: 1px solid black; text-align: center;">Carga nuclear, 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>

 

Este texto é adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/6-5-periodic-variations-in-element-properties">Openstax Chemistry 2e, Section 6.5: Periodic Variations in Element Properties.</a>

Radiação Atómica e Carga Nuclear Efetiva

Os elementos em grupos da tabela periódica exibem comportamento químico semelhante. Essa semelhança ocorre porque os membros de um grupo têm o mesmo número ...

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Periodic Properties of the Elements

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KEY TERMS AND CONCEPTS

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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

periodic classification of the elements

How Valence Electrons Shape the Periodic Table

The periodic table organizes elements by increasing atomic number, and it also shows why elements in the same group behave alike. As the electron configurations are placed into the table, a repeating pattern appears in the outer shells of atoms. These outer electrons are called valence electrons, and they play the biggest role in chemical reactions. They are the highest-energy electrons in an atom, so they are easier to lose or share than core electrons. Valence electrons also help determine...

ionization energy

Ionization Energy Trends in the Periodic Table

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

How Ionic Charge Changes Ion Size

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...

electron affinity

Electron Affinity Trends on the Periodic Table

Electron affinity is the energy change that happens when a gaseous atom gains an electron and becomes a negative ion, or anion. For many elements, this process releases energy, so the electron affinity value is negative. For some elements, energy must be added instead, so the value is positive. Electron affinity also applies when more than one electron is added. The second electron affinity is the energy change for adding an electron to an anion to make a 2− ion. Later electron affinities...

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Alkali Metals: Reactivity Trends

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: Trends and Reactions

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: Sources, Uses, and Reactions

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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Analyze and interpret data to explain the mechanisms and properties of the two main types of intramolecular (ionic and covalent) bonds.

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Periodic Trends in Atomic Size

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