#electron affinity

[materialsscienceandengineering]

ISOLDE reveals fundamental property of astatine, the rarest element on Earth

A team of researchers using the ISOLDE nuclear-physics facility at CERN has measured for the first time the so-called electron affinity of the chemical element astatine, the rarest naturally occurring element on Earth. The result, described in a paper just published in Nature Communications, is important for both fundamental and applied research. As well as giving access to hitherto unknown properties of this element and allowing theoretical models to be tested, the finding is of practical interest because astatine is a promising candidate for the creation of chemical compounds for cancer treatment by targeted alpha therapy.

The electron affinity is the energy released when an electron is added to a neutral atom in the gas phase to form a negative ion. It is one of the most fundamental properties of a chemical element. Together with the ionization energy, the energy it takes to remove an electron from the atom, it defines several other traits of an element, such as its electronegativity—the ability of the element to attract shared electrons in chemical bonds between atoms.

Although astatine was discovered in the 1940s, knowledge of its properties has mostly been based on theoretical calculations or on extrapolation from the properties of its relatives in the periodic table; astatine is a member of the halogen family, which includes chlorine and iodine. This is because astatine is scarce on Earth, and the tiny amounts of the element that can be produced in the lab prevent the use of traditional techniques to measure its properties. One notable exception was a previous measurement at ISOLDE of the element's ionization energy.

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[tjchemteam-blog]

Could you please explain in detail the concept and trends of the following? I had a hard time understanding in class and from the textbook: electronegativity, electron affinity, atomic radius, ionic radius, and ionization energy.

EDIT: oops typo about atomic radius trends

Electronegativity: measure of how tightly an element will hold onto an electron; the higher the value, the more tight the hold. Thus, the difference in electronegativity between two elements allows one to examine the nature of their bonding (whether it is ionic, polar covalent, or pure covalent). For example, the extremely large difference in electronegativity between Na and Cl, with Cl having the larger value, means that Cl will pull the electron cloud of the Na-Cl bond greatly towards itself, allows one to confirm that the bond in rock salt, NaCl, is ionic. SIDENOTE: The descending order of nonmetal electronegativities goes roughly FONClBrISCP, which kinda sounds like a funny word, so it’s easy to remember!

TREND: Electronegativity increases across a period, and decreases down a group. Fluorine is the most electronegative element. These trends do not include noble gases.

Electron Affinity (EA):The amount of energy released when the atom/molecule recieves an electron. NOTE: This means that when you see a negative EA value, it implies that the atom will absorb energy when taking in the electron. For instance, the 2nd EA value of oxygen is negative, which makes sense since an electron is being added to an already negative ion O-, which is unfavorable (making the ion more negative), thus raising the energy of the product O2-.

TREND: Electron affinity really does not follow the trends too well, but overall tends to increase going across a period. The values going down a group are rather variable. Frankly, a table of values is probably your best bet.

Atomic Radius: Measures the mean distance from the electron cloud to the nucleus. The main thing to know about the atomic radius is that the larger the value, the farther away electrons are from the nucleus and consequently less tightly held they are. Atomic radius actually explains pretty much all the other periodic trends. The larger the radius, the more easy it is to take an electron, and the lower the ionization energy and electronegativity values. The smaller the radius, the more tightly the nucleus holds onto electrons and would hold onto any incoming electrons, thus demonstrating the higher ionization energy and electronegativity.

TRENDS: Atomic radius decreases across a period, which makes sense as effective nuclear charge increases and thus attracts the electrons greatly. Atomic radius increases down a group, which also makes sense as we go up an electron shell and consequently the ends of the electron cloud move further from the nucleus. However, atomic radius doesn’t increase significantly across the transition metals because new electrons are being added to the d orbitals, which are on a lower shell than the s and p orbitals, and do not increase the size of the electron cloud.

Ionic Radius: Measure of the distance from the electron cloud to the nucleus in an ion. This works the same way atomic radius does: a higher effective nuclear charge will decrease the ionic radius. The important thing to know for ionic radius is that the ions of many elements will often have the same electron configuration. For example, the oxygen ion is most commonly -2, while the fluorine ion is most commonly -1. After gaining their respective electrons, these two ions have the same number of electrons. So, to determine which one has a smaller ionic radius, look at their nuclei. The fluorine nucleus has more protons than the oxygen nucleus does, so it will attract electrons more strongly and decrease the ionic radius to become smaller than oxygen.

Ionization energy: a measure of the energy needed to remove an electron from an atom. The removal of an electron to ionize, or make an ion, of an element will always require energy. This goes hand in hand with electronegativity, which should make sense; an atom that holds electrons tightly will require more energy to ionize.

TRENDS: Since ionization energy and electronegativity are so closely related, it makes sense that their trends are similar too: generally, ionization energy increases across a period and decreases down a group. However, there are some exceptions. In period 1, the ionization energy for boron is less than beryllium, and oxygen’s is less than nitrogen’s. This is because, in the case of beryllium, the removed electron will be in a p orbital, which is of higher energy than boron’s s orbital electron, meaning that less energy is neaded to remove it. In oxygen, the electron to be removed is found in a full orbital with 2 electrons, which repel each other and make it easier for the electron to be removed than nitrogen’s electron, which is the only one in its orbital and experiences less electrostatic repulsion.

This is another long answer, so feel free to message us with any questions!

[whatisthisthedarkages]

Periodic Trends

[marypoppinit]

does anybody understand chemistry bc i have not fucking clue what im doing

[pablochem5]

Periodic Trends: Electron Affinity & Electronegativity

Electron Affinity

  The electron affinity of the elements varies across the periodic table, though most trends are only loosely held.

The electron affinity of an atom or molecule is the propensity for that particle to gain an electron. This is an exothermic process.

There are trends in electron affinity across and down the periodic table of elements, though they are not universal. Electron affinity generally increases across a period in the periodic table and sometimes decreases down a group.

The chemical rationale for changes in electron affinity across the periodic table is the increased effective nuclear charge across a period and up a group.

  As with ionization energy, there are two rules that govern the periodic trends of electron affinities:

Electron affinity becomes less negative down a group.

As the principal quantum number increases, the size of the orbital increases and the affinity for the electron is less. The change is small and there are many exceptions.

Electron affinity decreases or increases across a period depending on electronic configuration.

This occurs because of the same subshell rule that governs ionization energies.

Example:

Since a half-filled "p" subshell is more stable, carbon has a greater affinity for an electron than nitrogen.

Obviously, the halogens, which are one electron away from a noble gas electron configuration, have high affinities for electrons:

(More negative energy = greater affinity)

*Fluorine's electron affinity is smaller than chlorine's because of the higher electron - electron repulsions in the smaller 2p orbital compared to the larger 3p orbital of chlorine.

Keep in mind that the electron affinity (just like IE) is a measure of how stable the products are with respect to the reactants. If the products are much more stable, a large amount of energy will be released during the process and EA will be a large negative number. At the other extreme, if the reactants are much more stable than the products, then it becomes very difficult to add an electron and the EA will be positive.

There are enough exceptions to the periodic trends in electron affinity that it is worthwhile to consider electron affinity of specific groups in the periodic table.

Halogens (group 7A, F to At)  Most negative EA values, addition of an e- leads to noble gas configuration, very favorable.

Group 5A (N to Bi)  filled shell discourages addition of an electron, EA values less negative than neighbors (groups 4A & 6A).

Alkaline Earths (group 2A, Be to Ba)  Filled s subshell discourages addition of an electron, EA values nearly zero.

Noble Gases (Group 8A, He to Rn)  Completely filled shell strongly discourages addition of an electron, EA values are positive.

Electronegativity

What is electronegativity?

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The higher the electronegativity of an atom, the greater its attraction for bonding electrons.

The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7.

Electronegativity and Ionization Energy

Electronegativity is related to ionization energy. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted on electrons by the nucleus.

Electronegativity and Periodic Table Trends

In an element group, the electronegativity decreases as atomic number increases, as a result of increased distance between the valence electron and nucleus (greater atomic radius). An example of an electropositive (i.e., low electronegativity) element is cesium; an example of a highly electronegative element is fluorine.

  Moving left to right across the periodic table, electronegativity increases.

Moving top to bottom down the periodic table, elctronegativity decreases.

Why this happens?

Electronegativity increases as you go across a period because, the next element along has an extra proton and an extra electron. The extra proton is in the nucleus, which means that the nucleus has a greater charge on it. The extra electron goes into the orbital with the lowest available energy, but because we are going across a period this means that the principle number of the orbital stays the same as for the last element. This means that the valence electrons do not experience any extra shielding from the nucleus, so they feel a greater effective nuclear charge and are more strongly attracted by the nucleus. This stronger attraction means the electrons are pulled closer to the nucleus, meaning the atom gets smaller (atomic radius decreases across a period). Thus, a shared pair of electrons will be more strongly attracted to the nucleus and so the atom withdraws more electron density. It is more electronegative.

Electronegativity decreases as you go down a group. This is because as you go down a group, the principal number of the valence orbital increases, meaning that there is an extra 'shell' of electrons between the valence electrons and the nucleus. This means that the valence electrons experience greater shielding from the nucleus. This factor is more important than the increased number of protons in the nucleus and the increased charge on the nucleus. So despite the extra protons the valence electrons are less strongly attracted by the nucleus, and the electrons are not held as close to the nucleus (atom radius increases down a group). Thus, a shared pair of electrons will be less strongly attracted to the nucleus, so the atom withdraws less electron density. It is less electronegative.

The D block

You may have noticed that the d block elements/transition metals do not seem to follow these trends quite as well as the s and p block elements, and there are indeed many exceptions to the trends in the d block. This is due to the valence electrons being in the d orbitals, these do not shield other electrons in the same way that s and p orbitals do. They do not shield electrons anywhere near as much as the s and p orbitals.

Additional facts to remember:

No electronegativity difference between two atoms leads to a pure non-polar covalent bond.

A small electronegativity difference leads to a polar covalent bond.

A large electronegativity difference leads to an ionic bond.