Periodic Trend: Electron Affinity: Videos & Practice Problems

Electron affinity (EA) is the energy released, typically measured in kilojoules, during the exothermic process of adding an electron to a gaseous atom or ion, resulting in a more negative species. A higher negative value of electron affinity indicates a stronger tendency to gain an electron. As we move across the periodic table towards the top right corner, electron affinity generally becomes more negative, signifying an increase in exothermicity. However, there are exceptions where elements have electron affinities that are zero or positive, indicating a reluctance to accept electrons. This reluctance is due to the stability provided by completely filled or half-filled orbitals, as seen in elements like beryllium, nitrogen, and noble gases. Elements like zinc, which have a completely filled d subshell, also exhibit this behavior. Understanding electron affinity is crucial as it influences how elements form chemical bonds and react with one another.

Electron Affinity

Electron affinity, denoted as \( e_a \), refers to the energy released when an electron is added to a gaseous atom or ion, typically measured in kilojoules. This process is exothermic, meaning it releases energy. For instance, when an electron is added to a neutral carbon atom, it transforms into a negatively charged ion, represented as \( \text{C}^- \) in the gaseous state. The release of energy during this process results in a negative value for electron affinity.

The magnitude of the electron affinity value indicates the extent of the exothermic reaction; the more negative the value, the more energy is released, signifying a stronger tendency for the atom to accept an electron. However, there are exceptions to this trend, particularly when examining the periodic table. If the electron affinity is equal to or greater than zero, it suggests that the element is less inclined to accept an electron, as doing so would disrupt its stability.

Understanding electron affinity is crucial for predicting how elements interact with electrons, which plays a significant role in chemical bonding and reactivity.

Electron Affinity Trends

Electron affinity is a key concept in understanding how atoms interact with electrons. It refers to the energy change that occurs when an electron is added to a neutral atom, and it is typically expressed in kilojoules per mole (kJ/mol). As we examine the periodic table, we observe that electron affinity generally becomes more exothermic, or more negative, as we move towards the top right corner. This trend indicates that elements in this region, such as fluorine, have a stronger tendency to attract electrons compared to those in the lower left, like cesium.

For instance, hydrogen has an electron affinity of -72.8 kJ/mol, while fluorine's is significantly more negative at -328 kJ/mol. This means fluorine has a greater desire to gain an electron than hydrogen. Conversely, cesium, with an electron affinity of -45.5 kJ/mol, shows a much weaker attraction for additional electrons.

Interestingly, some elements, particularly those in the bottom row of the periodic table, do not have reported electron affinity values. This is primarily due to their large size and instability, which complicates their ability to attract electrons. Additionally, certain elements like beryllium, nitrogen, and the noble gases exhibit electron affinities that are equal to or greater than 0. This indicates that these elements do not favor the addition of an electron, as they are already stable in their current state.

Other elements, such as manganese, zinc, cadmium, and mercury, also show electron affinities that are equal to or greater than 0, though the reasons for this behavior can vary. Notably, hydrogen fluoride (HF) is an anomaly, as its electron affinity does not fit neatly into the established trends, lacking a clear justification for its value being greater than or equal to 0.

Understanding these trends in electron affinity helps to explain the reactivity and stability of different elements, providing insight into their chemical behavior and interactions.

Electron Affinity Example

When considering which halogen has the most exothermic electron affinity, it's essential to focus on the elements in Group 7A of the periodic table, known as halogens. Among the options provided—sulfur, neon, nitrogen, astatine, and bromine—only astatine and bromine are halogens. To determine which of these has a more exothermic electron affinity, we need to understand that a more exothermic electron affinity corresponds to a more negative value.

As we move towards the top right corner of the periodic table, the electron affinity generally increases, indicating that elements become more exothermic in their electron affinity. In this context, bromine, being higher up in Group 7A compared to astatine, exhibits a more exothermic electron affinity. Therefore, bromine is the halogen with the most exothermic electron affinity, reflecting its ability to release more energy when gaining an electron.

Electron Affinity Exceptions

Electron affinity refers to the energy change that occurs when an electron is added to a neutral atom. However, certain elements exhibit exceptions to this trend due to their stable or symmetrical electron configurations. Elements with completely filled or half-filled orbitals tend to be less likely to accept additional electrons, as this would disrupt their stability.

For instance, Beryllium (Be) has a completely filled 2s orbital, resulting in an electron affinity of greater than or equal to 0, indicating a reluctance to gain an electron. Similarly, Nitrogen (N) possesses half-filled 2p orbitals, which are particularly stable; thus, it also resists accepting an electron. Manganese (Mn) showcases a similar behavior with its half-filled 3d orbitals, further illustrating this trend.

Moreover, Zinc (Zn) and elements below it in the periodic table have fully filled d orbitals, contributing to their stability and reluctance to gain electrons. These exceptions highlight the importance of electron configuration in determining an element's electron affinity.

In general, as one moves towards the top right corner of the periodic table, electron affinity tends to become more exothermic, meaning that the process of gaining an electron releases more energy. Understanding these exceptions is crucial for predicting the behavior of elements in chemical reactions.

Which of the following represents the third electron affinity of Si?

• A.
• B.
• C.
• D.

Answer: (C)

Determine which atom in the following set has the most exothermic electron affinity:N, O, C, B, Ne

• A. N
• B. O
• C. C
• D. B
• E. Ne

Answer: (B)

Rank the following elements in order of increasing electron affinity: Cs, Hg, F, S

• A. Hg < S < F < Cs
• B. Hg < F < S < Cs
• C. Cs < S < F < Hg
• D. Hg < Cs < S < F
• E. Cs < Hg < S < F

Answer: (D)

Here’s what students ask on this topic:

What is electron affinity?

Electron affinity refers to the amount of energy released when an electron is added to a neutral atom in the gaseous state to form a negative ion. It's a measure of the tendency of an atom to accept an electron. This property is important because it gives insight into the reactivity of different elements, particularly nonmetals, which generally have higher electron affinities.

When an electron is added to an atom, energy can either be released or absorbed. If energy is released, the electron affinity is positive, indicating that the atom readily accepts an electron. Conversely, if energy is absorbed, the electron affinity is considered negative, suggesting that the atom is less inclined to gain an electron.

Electron affinity can be influenced by several factors, including the atomic structure of the element and the effective nuclear charge—the net charge an electron experiences from the nucleus, taking into account the shielding effect of electrons in inner shells. Elements with a higher effective nuclear charge typically have greater electron affinities.

Why does electron affinity decrease down a group?

Electron affinity refers to the amount of energy released when an electron is added to a neutral atom in the gaseous state to form a negative ion. As you move down a group in the periodic table, the electron affinity generally decreases due to several factors.

Firstly, as you go down a group, the atomic size increases because additional electron shells are being added. This means that the added electron is further away from the nucleus, which lessens the attractive force exerted by the positively charged nucleus on the negative electron.

Secondly, there is an increase in electron shielding with each added electron shell. The inner shells of electrons act as a shield between the nucleus and the valence shell where the new electron is being added, reducing the effective nuclear charge felt by the added electron.

Because of these factors, the energy released when an electron is added to an atom lower in a group is typically less than that for an atom higher up in the same group, indicating a decrease in electron affinity.

Why does electron affinity increase across a period?

Electron affinity generally increases across a period on the periodic table due to the effective nuclear charge experienced by the electrons. As you move from left to right across a period, the number of protons in the nucleus increases, which in turn increases the nuclear charge. However, the additional electrons are being added to the same energy level and do not significantly increase the shielding effect. This means that the effective nuclear charge felt by the outermost electrons becomes stronger.

As a result, the attraction between the nucleus and the added electron becomes greater, making it more energetically favorable for the atom to gain an electron. This increased attraction leads to a higher electron affinity, meaning that atoms will release more energy when they gain an electron. However, there are exceptions to this trend, particularly when filled or half-filled subshells provide extra stability to the atom, which can affect its willingness to accept additional electrons.