Question

Select the element in each pair with the higher ionization energy. a. Br or \(I\) b. Mg or Sr c. Si or \(\overline{\mathrm{E}}\) d. I or Xe

Short answer

a. Br, b. Mg, c. Si, d. Xe

Step-by-step solution

Understand Ionization Energy

Ionization energy is the amount of energy required to remove an electron from an atom. Generally, ionization energy increases across a period from left to right and decreases down a group in the periodic table.

Compare elements in the same group

a. Br or I: Both Bromine (Br) and Iodine (I) are in Group 17, with Br being above I. Ionization energy decreases down the group, so Br has a higher ionization energy than I.

Compare elements in different groups

b. Mg or Sr: Magnesium (Mg) is in Group 2 and Period 3, while Strontium (Sr) is in Group 2 and Period 5. Ionization energy decreases down the group, so Mg has a higher ionization energy than Sr.

Compare elements in the same period

c. Si or E: Without the actual element _, it's unclear how to compare, so assuming E is in Group 14 with Si. Assuming all elements to the right generally possess higher ionization energies.

Compare elements in the same period and group

d. I or Xe: Iodine (I) is in Group 17, whereas Xenon (Xe) is in Group 18. Typically, elements in Group 18 have very high ionization energies as they're noble gases. Thus, Xe has a higher ionization energy than I.

Key concepts

Ionization Energy Trend

Ionization energy is the energy required to remove an electron from an atom in the gaseous state. The trend of ionization energy in the periodic table is quite predictable. As you move from left to right across a period, ionization energy increases. This happens because the nuclear charge increases, pulling the electrons closer to the nucleus. As a result, more energy is needed to remove an electron. Conversely, as you move down a group, ionization energy decreases. This is because the outer electrons are farther from the nucleus and experience more shielding from inner-shell electrons. Hence, less energy is needed to remove an electron.

Periodic Table Groups

The periodic table is organized into vertical columns known as groups or families. Each group contains elements with similar chemical properties. One reason for this is that elements in the same group have the same number of electrons in their outermost shell. For example, Group 1 elements, also known as alkali metals, all have one electron in their outer shell. This similarity in electronic configuration leads to similarities in chemical behavior. Generally, the ionization energy decreases as you go down a group because the atoms get larger and the outer electrons are increasingly distant from the nucleus.

Periodic Table Periods

The periodic table is also organized into horizontal rows known as periods. As you move from left to right across a period, the number of protons and electrons in an atom increases. This leads to stronger electrostatic forces between the positively charged nucleus and the negatively charged electrons. Consequently, the ionization energy increases across a period. For example, nitrogen (N) has a higher ionization energy than lithium (Li) because nitrogen is further to the right in the same period, resulting in a higher nuclear charge drawing the electrons in more tightly.

Group 17 Elements

Group 17 elements are known as halogens. These elements include fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). Halogens have seven electrons in their outermost shell, making them extremely reactive as they seek to gain one more electron to achieve a stable octet configuration. The ionization energy in Group 17 decreases as you move down the group. For instance, bromine (Br) has a higher ionization energy than iodine (I), as Br is higher up in the group. Halogens have high electronegativities and are typically very reactive non-metals.

Group 18 Elements

Group 18 elements are known as noble gases. These elements include helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). Noble gases are characterized by having a complete outer electron shell, which makes them very stable and unreactive. Because their outer shell is full, they do not easily lose or gain electrons. This results in them having very high ionization energies. For example, xenon (Xe) has a higher ionization energy than iodine (I) despite being in the same period because Xe is a noble gas with a complete electron shell, thereby requiring more energy to remove an electron.