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

Das chemische Verhalten von Atomen und Ionen wird stark davon beeinflusst, wie einfach oder schwierig es ist, ihre Elektronen zu entfernen, insbesondere die äußersten Elektronen, die an der chemischen Verbindung beteiligt sind. Die zum Entfernen eines Elektrons aus einem gasförmigen Atom in seinem Grundzustand erforderliche Energie wird die erste Ionisationsenergie genannt und wird in kJ/mol angegeben. Die zum Entfernen des nächsten Elektrons erforderliche Energie wird die zweite Ionisationsenergie genannt, und so weiter.Spaltenabwärts, nehmen die Ionisationsenergien ab. Erinnern Sie sich, dass die höchste Hauptquantenzahl der Valenzelektronen in der Spalte abnimmt, was zu größeren Atomgrößen führt. Je weiter also die äußersten Elektronen entfernt sind, desto leichter sind sie zu entfernen.Für Hauptgruppenelemente wird die Ionisationsenergie im Laufe der Periode steigen. Der Grund liegt in der steigenden Ordnungszahl, wobei Valenzelektronen eine höhere effektive Kernladung haben, was die Entfernung der äußersten Elektronen erschwert. Dies erklärt, warum Chlor eine höhere Ionisationsenergie als zum Beispiel Natrium hat.Im Allgemeinen ist die Ionisierungenergie bei einem Alkalimetall minimal und steigt mit jedem Edelgas bis zum Maximum. Übergangsmetalle zeigen über die Periode einen geringen Anstieg in der Ionisationsenergie, und die f Block-Elemente zeigen eine noch geringere Veränderung. Aber es gibt einige Ausnahmen zu beachten.Bor hat eine geringere Ionisationsenergie als Beryllium, auch wenn es im Periodensystem weiter rechts liegt. Beryllium hat 2s Elektronen niedrigerer Energie, wohingegen Bor eine höhere Energie hat, das 2p Elektron, wodurch seine Entfernung energetisch günstiger wird. Eine weitere Ausnahme ist Sauerstoff, der eine niedrigere Erst-Ionisationsenergie als Stickstoff hat.Verglichen zu Stickstoff hat Sauerstoff vier p Elektronen, und das Entfernen eines Elektrons eliminiert die Elektron-Elektron-Abstoßung. Dadurch wird weniger Energie für die Ionisierung benötigt. Diese Ausnahmen sind auch in den Folgeperioden zu beobachten.Elektronenentfernung aus Kationen ist schwieriger als aus neutralen Atomen. Im Allgemeinen steigen die aufeinanderfolgenden Ionisationsenergien der Elemente. Berücksichtigen Sie Kalium.Die zweite Ionisationsenergie ist deutlich höher, da es die Entfernung eines Kernelektrons aus einem Ion mit einer Edelgaskonfiguration beinhaltet. Ähnlich ist es beim Calcium, wo ein hoher Anstieg zu verzeichnen ist, von der zweiten zur dritten Ionisationsenergie, wenn ein Kernelektron aus einem Kation mit einer Edelgaskonfiguration entfernt wird.

ionisationsenergie

Die Energiemenge die erforderlich ist, um das am wenigsten stark gebundene Elektron aus einem gasförmigen Atom im Grundzustand zu entfernen, wird als erste Ionisierungsenergie (IE1) bezeichnet. Die erste Ionisierungsenergie für ein Element, X, ist die Energie, die erforderlich ist, um ein Kation mit der Ladung 1+ zu bilden: <img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1049/1049_Equations-01.png" style="height: 40px;"> Die Energie, die erforderlich ist, um das am zweitleichtesten gebundene Elektron zu entfernen, wird als zweite Ionisierungsenergie (IE2) bezeichnet. <img alt="Eq2" src="https://www.jove.com/files/core/CDNSource/lessons/1049/1049_Equations-02.png" style="height: 40px;"> Die zum Entfernen des dritten Elektrons erforderliche Energie ist die dritte Ionisierungsenergie und so weiter. Um Elektronen aus Atomen oder Ionen zu entfernen, ist immer Energie erforderlich, daher sind Ionisierungsprozesse endotherm und die IE-Werte sind immer positiv. Bei größeren Atomen befindet sich das am wenigsten gebundene Elektron weiter vom Kern entfernt und lässt sich daher leichter entfernen. Mit zunehmender Größe (Atomradius) sollte daher die Ionisierungsenergie abnehmen.  Innerhalb einer Periode nimmt der IE1-Wert im Allgemeinen mit zunehmendem Z zu. In einer Gruppe nimmt der IE1-Wert im Allgemeinen mit zunehmendem Z ab. Es gibt jedoch einige systematische Abweichungen von diesem Trend. Zu beachten ist, dass die Ionisierungsenergie von Bor (Ordnungszahl 5) geringer ist als die von Beryllium (Ordnungszahl 4), obwohl die Kernladung von Bor um ein Proton größer ist. Dies kann dadurch erklärt werden, dass die Energie der Unterschalen aufgrund der Durchdringung und Abschirmung mit zunehmendem l zunimmt. Innerhalb jeder Schale haben die s-Elektronen eine niedrigere Energie als die p-Elektronen. Das bedeutet, dass ein s-Elektron schwerer aus einem Atom zu entfernen ist als ein p-Elektron der selben Schale. Das bei der Ionisierung von Beryllium ([He]2s2) entfernte Elektron ist ein s-Elektron, während das bei der Ionisierung von Bor ([He]2s22p1) entfernte Elektron ein p-Elektron ist; Dies führt zu einer niedrigeren ersten Ionisierungsenergie für Bor, obwohl seine Kernladung um ein Proton größer ist. Daher sehen wir jedes Mal, wenn eine neue Unterschale beginnt, eine kleine Abweichung vom vorhergesagten Trend. Eine weitere Abweichung tritt auf, wenn die Orbitale mehr als zur Hälfte gefüllt sind. Die erste Ionisierungsenergie für Sauerstoff ist etwas geringer als die für Stickstoff, trotz der Tendenz der Erhöhung, der IE1-Werte entlang der Periode. Bei Sauerstoff wird durch das Entfernen eines Elektrons die Elektron-Elektron-Abstoßung veringert, die durch die Paarung der Elektronen im 2p-Orbital verursacht wird und es entsteht ein (energetisch begünstigtes) halbgefülltes Orbital. Analoge Veränderungen treten auch in den darauf folgenden Perioden auf. Das Entfernen eines Elektrons von einem Kation ist aufgrund der größeren elektrostatischen Anziehungskraft des Kations schwieriger als das Entfernen eines Elektrons von einem neutralen Atom. Ebenso ist das Entfernen eines Elektrons von einem Kation mit einer höheren positiven Ladung schwieriger als das Entfernen eines Elektrons von einem Ion mit einer niedrigeren Ladung. Somit nehmen aufeinanderfolgende Ionisierungsenergien für ein Element immer zu. Wie in Tabelle 1 zu sehen ist, gibt es an einer bestimmten Stelle einen starken Anstieg der Ionisierungsenergien bei jedem Element. Dieser Sprung entspricht dem Entfernen der Kernelektronen, die schwerer vom Atom zu trennen sind als die Valenzelektronen. Sc und Ga haben beispielsweise beide drei Valenzelektronen, sodass in diesem Fall der starke Anstieg der Ionisierungsenergie nach der dritten Ionisierung erfolgt. Tabelle 1: Aufeinanderfolgende Ionisierungsenergien für ausgewählte Elemente (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; ">Element</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">IE1</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">IE2</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">IE3</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">IE4</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">IE5</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">IE6</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">IE7</td> </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> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">589.8</td> <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> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">13315.0</td> </tr> <tr> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">Ga</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">578.8</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">1979.4</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">2964.6</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">6180</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">8298.7</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">10873.9</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">13594.8</td> </tr> <tr> <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> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">1537.5</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">3302.1</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">4410.6</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">9021.4</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">Nicht verfügbar</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">Nicht verfügbar</td> </tr> <tr> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">As</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">944.5</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">1793.6</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">2735.5</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">4836.8</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">6042.9</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">12311.5</td> <td style="color: black; border: 1px solid black; text-align: center; padding: 10px; ">Nicht verfügbar</td> </tr> </tbody> </table> Dieser Text wurde adaptiert von <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>

Ionisationsenergie

Die Energiemenge, die erforderlich ist, um das am lockersten gebundene Elektron aus einem gasförmigen Atom in seinem Grundzustand zu entfernen, wird als erste Ionisationsenergie bezeichnet ...

Education

Lehre

JoVE Core

Chemie

Eins

Periodic Properties of the Elements

Periodische Eigenschaften der Elemente

KEY TERMS AND CONCEPTS

KERNBEGRIFFE UND KONZEPTE

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...

halogens

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...

noble gases

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...

key terms and concepts

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 trends of the elements

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Periodic Table Trends

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Periodic Table Trends: Atomic Radii, Ionization, and Electronegativity

Scientific Laws and Theories

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

Chemistry is an empirical science. Scientists often pose questions to understand the chemistry in everyday life and seek answers to these questions. To ...

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Nuclear chemistry is the study of reactions that involve changes in nuclear structure. The nucleus of an atom is composed of protons and, except for ...

Types of Radioactivity

The most common types of radioactivity are &#945; decay, &#946; decay, &#947; decay, neutron emission, and electron capture.
Alpha (&#945;) decay is the ...

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Protons and neutrons, collectively called nucleons, are packed together tightly in a nucleus. With a radius of about 10&#8722;15 meters, a nucleus is ...

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The difference between the calculated and experimentally measured masses is known as the mass defect of the atom. In the case of helium-4, the mass defect ...

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Radioactivity is a spontaneous disintegration of an unstable nuclide and is a random process, as all the nuclei in the sample do not decay simultaneously. ...

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Many heavier elements with smaller binding energies per nucleon can decompose into more stable elements that have intermediate mass numbers and larger ...

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Atomic Orbitals

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The Pauli Exclusion Principle

The arrangement of electrons in the orbitals of an atom is called its electron configuration. We describe an electron configuration with a symbol that ...

The Energies of Atomic Orbitals

In an atom, the negatively charged electrons are attracted to the positively charged nucleus. In a multielectron atom, electron-electron repulsions are ...

The Aufbau Principle and Hund's Rule

To determine the electron configuration for any particular atom, we can build the structures in the order of atomic numbers. Beginning with hydrogen, and ...

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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 ...

Use models to explain how electrons are distributed in atoms.

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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

As early chemists discovered more elements, they realized that various elements could be grouped by their similar chemical behaviors. One such grouping ...

Atomic Mass

Atoms &mdash; and the protons, neutrons, and electrons that compose them &mdash; are extremely small. For example, a carbon atom weighs less than 2 ...

Molar Mass

The identity of a substance is defined not only by the types of atoms or ions it contains but by the quantity of each type of atom or ion. For example, ...

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PS.Chm.2.1

Atomic Radii and Effective Nuclear Charge

The elements in groups of the periodic table exhibit similar chemical behavior. This similarity occurs because the members of a group have the same number ...

Ionic Radii

Ionic radius is the measure used to describe the size of an ion. A cation always has fewer electrons and the same number of protons as the parent atom; it ...

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

Chemical bonding theories were pioneered by American chemist Gilbert&#xA0;N. Lewis. He developed a model called the Lewis model to explain the type and ...

Lewis Symbols and the Octet Rule

Chemical bonds are complex interactions between two or more atoms or ions, which reduce the potential energy of the molecule. Gilbert N. Lewis developed a ...

Ionic Bonding and Electron Transfer

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The Born-Haber Cycle

Lattice Energy&#xA0;

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

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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 ...

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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