Even though adding an electron would generate an ion (anion), ionization energy is defined as the energy needed to remove an electron.

As you ionize atoms, you generate charged ionic species. These charges will resist further ionization (cations will more strongly attract the electrons you are trying to pull away). Noble gas configurations are particularly stable, so you would expect a large increase in needed energy to ionize away from this state.

Cesium, tungsten, platinum, and radon are all in the same row (period) of the periodic table.

The ionization energy of an atom, defined as the energy required to remove an electron from an atom’s orbit, increases as one move from left to right across the periodic table. This is primarily because the atoms decrease in atomic size, meaning that the outermost electrons that would be removed to create an ion are closer to the nucleus and require more work to remove. Cesium has the largest atomic radius and attains an octet if an electron is removed, making it more stable as an ion, meaning that it will be fairly easy to remove an electron from cesium compared to the other options.

Fluorine, chlorine, bromine, and iodine are all halogens in the same group of the periodic table.

The ionization energy of an atom, defined as the energy required to remove an electron from an atom’s orbit, decreases as one moves down a group of the periodic table. This is primarily because the atoms increase in size as one moves down the group, thus decreasing the force between the nucleus and an electron in an outermost orbit. It requires less energy to remove an electron that is farther from the protons, than to remove one that is held closer to the nucleus.

Ionization energy describes the energy required to remove a single electron from an atom. The periodic table trend for ionization energy is to increase from left to right or a period, and decrease from top to bottom of a group. Typically, elements that have seven valence electrons, one electron away from having a complete valence shell, have very high ionization energies. These elements are found adjacent to noble gases on the periodic table: the halogens. Chlorine would thus have a very high ionization energy.

Alkali metals have the lowest ionization energies, as removing a single electron will give them a full octet configuration.

Noble gasses, like krypton, have a full octet in their valence shells. Their stability makes it very hard to alter their valence configuration. First ionization energy increases from left to right across the periodic table.

Bromine, which is right next to krypton, also has a high first ionization energy because it would much rather capture an electron than release one. As you move down the table, the energy for that first ionization gets lower, but as you move across from left to right it gets higher. Transition elements don't follow the trend exactly, but the energy required for them is not higher than halogens or noble gasses.

Alkali metals, like lithium, have the lowest first ionization energies because they obtain a full octet by releasing an electron.

Ionization energy refers to the amount of energy necessary to remove a single electron from an atom or ion in the gas state, resulting in an increase in charge.

To find our answer, we need to know that ionization energy increases as we move left to right across a period and decreases as we travel down a group. Alkali metals, on the left side of the table, gain stability with the removal of a single electron, while halogens lose stability. Atoms or ions with a electron configurations similar to those of the noble gases are very stable. Due to this stability, larger amounts of energy are required to alter the electron configuration.

When a halogen gains an electron, it achieves a stable valence octet; thus, the fluoride and bromide ions are incredibly stable and will have very high first ionization energies. Bromine is lower on the periodic table than fluorine, so its electrons are separated from its positively-charged nucleus by more energy shells. This effect, called shielding, leads to a smaller amount of energy required to remove one electron from bromine. , therefore, is our solution.

Ionization energy is the energy required by an element to remove a valence electron and gain a positive charge. Because alkali metals and alkaline earth metals have one and two valence electrons, respectively, their first ionization energies are both relatively low; shedding a valence electron and absorbing a positive charge actually stabilizes these elements.

Alkaline earth metals can easily release their second valence electron, and have relatively low second ionization energies. Removing this second electron gives these elements a noble gas electronic configuration. This is why the ions of alkaline earth metals are most commonly Mg²⁺ and Ca²⁺. In comparison, alkali metals cannot easily shed a second valence electron because it would require removing an electron from an already-filled valence shell. Because of this, alkali metals have extremely HIGH second ionization energies in comparison to alkaline earth metals.

Of the answer choices, potassium would have the highest second ionization energy because it is an alkali metal, rather than an alkaline earth metal.

Rubidium, ruthenium, tin, and xenon are all in the same row (period) of the periodic table.

Metallic character decreases as one moves from left to right across a period. This is generally due to a reduction in the number of electrons in the outermost shell being from the s orbital. Increased numbers of protons reduce the atomic radius, making it harder for outer shells to have electron motility. Metals will generally have larger radii, allowing for greater electron freedom, ionization, and conductivity. Remember that metallic characteristics include malleability, ductility, luster, and shininess.

Lithium, potassium, rubidium, and francium are all alkali metals in the same group of the periodic table.

Metallic character increases as one moves down each group the periodic table. This is generally due to an increase in the distance between the electrons and the nuclear protons, making it easier for an electron to be lost to a bond or to move freely during electric conduction. Remember that metallic characteristics include malleability, ductility, luster, and shininess.