Name: ionization energy trends in the periodic table
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 3 min 12 s
Channel: JoVE

K12 Standards

north-carolina-k-12-science-standard-course-of-study

NC K-12 Science Standard Course of Study

9-12

Chemistry

Matter and its Interactions

PS.Chm.2.2

Ionization Energy

Le comportement chimique des atomes et des ions est grandement affecté par la facilité ou la difficulté d'éliminer leurs électrons, en particulier les électrons les plus externes qui participent aux formations de liaisons chimiques. L'énergie nécessaire pour éliminer un électron d'un atome gazeux dans son état fondamental est appelée la première énergie d'ionisation qio est est calculée en kilojoules par mole. L'énergie nécessaire pour enlever l'électron suivant est appelée la deuxième énergie d'ionisation, et ainsi de suite.En descendant une colonne, les énergies d'ionisation diminuent. Rappelons que le nombre quantique principal le plus élevé d'électrons de valence augmente dans la colonne conduisant à des tailles atomiques plus grandes. Ainsi, plus les électrons les plus éloignés sont, plus ils sont faciles à enlever.Pour les éléments du groupe principal, l'énergie d'ionisation augmente au cours de la période. La raison est liée à l'augmentation du nombre d'atomes, où les électrons de valence subissent une charge nucléaire plus efficace, ce qui rend plus difficile l'élimination des électrons les plus éloignés. Ceci explique pourquoi le chlore a une énergie d'ionisation plus élevée que le sodium, par exemple.Généralement, l'énergie d'ionisation est un minimum pour un métal alcalin et s'élève à un pic avec chaque gaz noble. Les métaux de transition affichent une légère augmentation de l'énergie d'ionisation au cours de la période, et les éléments du bloc f montrent un changement encore plus faible. Mais il y a quelques exceptions à considérer.Le bore a une énergie d'ionisation plus petite que le béryllium, même s'il est plus à droite sur le tableau périodique. Le béryllium a des électrons 2s d'énergie plus faibles, tandis que le bore a un électron 2p d'énergie plus élevée, ce qui rend son élimination énergétiquement plus favorable. Une autre exception est l'oxygène, qui a une énergie de première ionisation inférieure à l'azote.Comparé à l'azote, l'oxygène a quatre p-électrons, et supprimer d'un électron élimine l'électron-répulsion électron. Ainsi, moins d'énergie est nécessaire pour l'ionisation. Ces exceptions sont également observées dans les périodes suivantes.L'élimination des électrons des cations est plus difficile que celle des atomes neutres. Généralement, les énergies d'ionisation successives augmentent pour les éléments. Considérez le potassium.La seconde énergie d'ionisation est significativement plus élevée, car elle implique l'élimination d'un électron central d'un ion avec une configuration de gaz noble. De même, pour le calcium, il y a une augmentation élevée de la deuxième à la troisième énergie d'ionisation car un électron central est retiré d'un cation avec une configuration de gaz noble.

Énergie d'ionisation

La quantité d'énergie nécessaire pour enlever l'électron le moins lié d'un atome gazeux dans son état fondamental est appelée son énergie de première ionisation (Ei1). L'énergie de première ionisation pour un élément, X, est l'énergie requise pour former un cation avec une charge de 1+ : 
 
 <img alt="Eq1" src="https://www.jovecore.com/CDNSource/lessons/1049/1049_Equations-01.png" style="height: 40px;" />
 
 L'énergie nécessaire pour enlever le deuxième électron le moins lié est appelée l'énergie de deuxième ionisation (Ei2). 
 
 <img alt="Eq2" src="https://www.jovecore.com/CDNSource/lessons/1049/1049_Equations-02.png" style="height: 40px;" />
 
 L'énergie nécessaire pour enlever le troisième électron est l'énergie de troisième ionisation, et ainsi de suite. L'énergie est toujours nécessaire pour arracher les électrons des atomes ou des ions, de sorte que les processus d'ionisation sont endothermiques et les valeurs d'Ei sont toujours positives. Pour les atomes plus grands, l'électron le moins lié est situé plus loin du noyau et il est donc plus facile à enlever. Ainsi, à mesure que la taille (rayon atomique) augmente, l'énergie d'ionisation devrait diminuer.  
 
 Dans une période, Ei1 augmente généralement avec le Z croissant. En descendant dans un groupe, la valeur d'Ei1 diminue généralement avec le Z croissant. Il y a toutefois quelques écarts systématiques par rapport à cette tendance. Il convient de noter que l'énergie d'ionisation du bore (numéro atomique 5) est inférieure à celle du béryllium (numéro atomique 4), même si la charge nucléaire du bore est plus importante d'un proton. Ceci peut être expliqué par le fait que l'énergie des sous-couches augmente à mesure que l augmente, en raison de la pénétration et de l'effet d'écran. Dans n'importe quelle couche, les électrons s sont plus faibles en énergie que les électrons p. Cela signifie qu'un électron s est plus difficile à arracher d'un atome qu'un électron p sur la même couche. L'électron arraché lors de l'ionisation du béryllium ([He]2s2) est un électron s, alors que l'électron arraché lors de l'ionisation du bore ([He]2s22p1) est un électron p, ce qui se traduit par une énergie de première ionisation plus faible pour le bore, même si sa charge nucléaire est plus grande d'un proton. Ainsi, par rapport à la tendance prévue, nous voyons un petit écart se produire chaque fois qu'une nouvelle sous-couche commence. 
 
 Un autre écart se produit lorsque les orbitales sont remplies à plus de la moitié. L'énergie de première ionisation de l'oxygène est légèrement inférieure à celle de l'azote, malgré la tendance à l'élévation des valeurs de Ei1 sur une période donnée. Dans le cas de l'oxygène, l'arrachement d'un électron éliminera la répulsion électron-électron causée par l'appariement des électrons dans l'orbitale 2p et le résultat sera une orbitale à moitié remplie (ce qui est énergétiquement favorable). Des changements analogues se produisent dans les périodes qui suivent. 
 
 Il est plus difficile d'arracher un électron à un cation que d'arracher un électron à un atome neutre en raison de la plus grande attraction électrostatique vers le cation. De même, il est plus difficile d'arracher un électron à un cation ayant une charge positive plus élevée que d'arracher un électron à un ion ayant une charge plus faible. Ainsi, les énergies d'ionisation successives pour un élément augmentent toujours. Comme le montre le tableau 1, il y a une forte augmentation des énergies d'ionisation pour chaque élément. Ce saut correspond à l'arrachement des électrons de cœur, qui sont plus difficiles à arracher que les électrons de valence. Par exemple, Sc et Ga ont tous les deux trois électrons de valence, de sorte que l'augmentation rapide de l'énergie d'ionisation se produit après la troisième ionisation. 
 
 Tableau 1 : énergies d'ionisation successives pour quelques éléments sélectionnés (kJ/mol) 
 
 <table style="margin-left: auto; margin-right: auto;">
 <tbody style="border: 1px solid black;">
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">Élément </td>
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 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">Ei 2</td>
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 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">Ei 4</td>
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 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">K </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">418,8 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">3051,8 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">4419,6 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">5876,9 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">7975,5 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">9590,6 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">11343 </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">Ca </td>
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 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">1145,4 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">4912,4 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">6490,6 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">8153,0 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">10495,7 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">12272,9 </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">Sc </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">633,1 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">1235,0 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">2388,7 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">7090,6 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">8842,9 </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">10679,0 </td>
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 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">Ga </td>
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 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">Ge </td>
 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">762,2 </td>
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 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">As </td>
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 <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">Non disponible</td>
 </tr>
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 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>

La quantité d’énergie nécessaire pour éliminer l’électron le plus faiblement lié d’un atome gazeux dans son état fondamental est appelée sa première énergie d’ionisation...

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Periodic Properties of the Elements

Propriétés périodiques des éléments

KEY TERMS AND CONCEPTS

TERMES ET CONCEPTS CLÉS

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 atoms have lower ionization energy because their outermost electrons are farther from the nucleus and easier to remove.

Across a period, first ionization energy generally increases as atomic number increases. Down a group, first ionization energy generally decreases as atomic number increases. The trend is useful, but there are some clear exceptions.

One exception appears when a new subshell begins. Boron has a lower first ionization energy than beryllium, even though boron has one more proton. This happens because s electrons are lower in energy than p electrons in the same shell, so the s electron in beryllium is harder to remove than the p electron in boron. The effect comes from penetration and shielding.

Another exception appears when a p subshell becomes more than half filled. Oxygen has a slightly lower first ionization energy than nitrogen. Removing one electron from oxygen reduces electron-electron repulsion in the 2p orbital and gives a more stable half-filled orbital.

Successive ionization energies always increase. It is harder to remove an electron from a cation than from a neutral atom because the positive charge pulls more strongly on the remaining electrons. It is even harder to remove an electron from a cation with a higher positive charge.

The large jump in successive ionization energies shows when core electrons begin to be removed. Core electrons are harder to remove than valence electrons. For example, scandium and gallium both have three valence electrons, so the sharp rise in ionization energy comes after the third ionization. The table of values for K, Ca, Sc, Ga, Ge, and As shows this pattern clearly.

Ionization Energy Trends in the Periodic Table

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 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

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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...

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...

alkali metals

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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Explore first and successive ionization energy, periodic trends, and key exceptions such as boron, oxygen, and the jump to core electrons.

Ionization Energy Trends in the Periodic Table | JoVE K12

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Periodic Table Trends: Atomic Radii, Ionization, and Electronegativity

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Scientific Laws

In science, a law is defined as a concise, verbal or mathematical, statement that summarizes a vast number of experimental observations. ...

The Scientific Method

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Alpha (&#945;) decay is the ...

Nuclear Stability

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Electron Configuration of Multielectron Atoms

The alkali metal sodium (atomic number 11) has one more electron than the neon atom. This electron must go into the lowest-energy subshell available, the ...

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Periodic Classification of the Elements

The periodic table arranges atoms based on increasing atomic number so that elements with the same chemical properties recur periodically. When their ...

The Periodic Table

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Molar Mass

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Atomic Radii and Effective Nuclear Charge

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Ionic Radii

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The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy ...

Electron Affinity

The electron affinity (EA) is the energy change for adding an electron to a gaseous atom to form an anion (negative ion).



This process can be either ...

Construct an explanation to infer the atomic size, reactivity, electronegativity, and ionization energy of an element based on its position in the Periodic Table.

Types of Chemical Bonds

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The Born-Haber Cycle

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Trends in Lattice Energy: Ion Size and Charge

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a compound is a measure ...

Covalent Bonding and Lewis Structures

Compared to ionic bonds, which results from the transfer of electrons between metallic and nonmetallic atoms, covalent bonds result from the mutual ...

Electronegativity

Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity.&#xA0;

Electronegativity values of ...

bond-polarity-dipole-moment-and-percent-ionic-character

Bond Polarity, Dipole Moment, and Percent Ionic Character

Bond Polarity

The absolute value of the difference in electronegativity (&Delta;EN) of two bonded atoms provides a rough measure of the polarity to be ...

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

PS.Chm.3.1

Analyze the structure of atoms and isotopes.

Ionization Energy Trends in the Periodic Table

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