Question

Answer the following questions based on the given electron configurations and identify the elements. a. Arrange these atoms in order of increasing size: \([\mathrm{Kr}] 5 s^{2} 4 d^{10} 5 p^{6} ;[\mathrm{Kr}] 5 s^{2} 4 d^{10} 5 p^{1} ;[\mathrm{Kr}] 5 s^{2} 4 d^{10} 5 p^{3}\) b. Arrange these atoms in order of decreasing first ionization energy: \([\mathrm{Ne}] 3 s^{2} 3 p^{5} ;[\mathrm{Ar}] 4 s^{2} 3 d^{10} 4 p^{3} ;[\mathrm{Ar}] 4 s^{2} 3 d^{10} 4 p^{5}\)

Short answer

The short answer based on the step-by-step solution is: a. The order of increasing atomic size for the given electron configurations is Rubidium (Rb) < Yttrium (Y) < Molybdenum (Mo). b. The order of decreasing first ionization energy for the given electron configurations is Chlorine (Cl) > Bromine (Br) > Arsenic (As).

Step-by-step solution

Identify the elements

First, let's identify the elements in the given electron configurations: 1. [Kr] 5s² 4d¹⁰ 5p⁶: The last noble gas in the electron configuration is Krypton (Kr), and the configuration corresponds to a fully-filled 5s, 4d, and 5p subshells. The atomic number of Kr is 36, and there are a total of 6 more electrons after Kr, so this element has an atomic number of 42 - which corresponds to Mo (Molybdenum). 2. [Kr] 5s² 4d¹⁰ 5p¹: Similar to the first configuration, the atomic number of Kr is 36, and there is 1 additional electron, so this configuration corresponds to an atomic number of 37, which is Rb (Rubidium). 3. [Kr] 5s² 4d¹⁰ 5p³: Here, the atomic number of Kr is 36, with 3 additional electrons, giving a total atomic number of 39, which is Y (Yttrium). 4. [Ne] 3s² 3p⁵: The electron configuration ends with a Neon (Ne) core, with an atomic number of 10. There are 7 more electrons, giving a total atomic number of 17, which is Cl (Chlorine). 5. [Ar] 4s² 3d¹⁰ 4p³: This configuration ends with an Argon (Ar) core, with an atomic number of 18. There are 15 more electrons in the configuration, giving a total atomic number of 33, which is As (Arsenic). 6. [Ar] 4s² 3d¹⁰ 4p⁵: The electron configuration ends with an Argon (Ar) core, with an atomic number of 18. There are 17 more electrons in the configuration, giving a total atomic number of 35, which is Br (Bromine). Now that we've identified the elements, we can arrange them according to atomic size and first ionization energy.

Part a. Arrange atoms in order of increasing size

Atomic size increases from top to bottom within a group and decreases from left to right across a period in the periodic table. The given elements, Rubidium (Rb), Molybdenum (Mo), and Yttrium (Y) are in the same period with Rb being in Group 1, Y in Group 3, and Mo in Group 6. Following the periodic trends, the size would decrease from left to right. Therefore, the order of increasing atomic size is Rb < Y < Mo.

Part b. Arrange atoms in order of decreasing first ionization energy

The first ionization energy increases from left to right across a period and decreases from top to bottom within a group in the periodic table. The given elements, Chlorine (Cl), Arsenic (As), and Bromine (Br) are in the same group (Group 17). Chlorine is in Period 3, while Arsenic and Bromine are in Period 4. Following the periodic trends, the ionization energy would decrease from top to bottom in a group. As they are in the same row, ionization energy would be similar, but As would have slightly lower ionization energy due to electron repulsion in 3d orbitals. Therefore, the order of decreasing first ionization energy is Cl > Br > As.

Key concepts

Periodic Trends

Understanding the behavior of elements in the periodic table is grounded on the recognition of patterns known as periodic trends. These trends explain the systematic variations in the properties of elements, including atomic size, ionization energy, and electron affinity among others. As we move across a period (left to right) in the periodic table, we typically see an increase in electronegativity and ionization energy, whereas atomic radius decreases. Conversely, moving down a group (top to bottom), elements tend to increase in atomic size due to an additional electron shell, while their ionization energy and electronegativity decrease.

This is because of two main factors: the increase in positive charge of the nucleus, which pulls electrons closer, and the increase in the number of electron shells, which can shield the outermost electrons from the pull of the nucleus. These patterns are crucial for predicting and explaining the chemical behavior of elements, as seen in the arrangement of atoms based on their size and ionization energy.

Atomic Size

The term atomic size refers to the distance from the nucleus of an atom to its outermost shell of electrons. Atomic size is a fundamental concept in understanding chemical bonding and reaction patterns.

The trend across the periodic table shows that atomic size decreases from left to right within a period. This is due to the increasing nuclear charge which draws the electron cloud inwards, making the atoms smaller. However, when moving down a group, the atomic size increases because each successive element has an additional electron shell, providing greater distance between the nucleus and the outermost electrons.

In the discussion of atomic configuration, elements with larger atomic sizes generally have lower first ionization energies since the outermost electrons are further from the nucleus and are easier to remove. The relationship between atomic size and its position on the periodic table is vital for predicting how atoms will interact with one another.

First Ionization Energy

The first ionization energy is defined as the energy required to remove the most loosely bound electron from a gaseous atom in its ground state. It is a measure of the tendency of an atom to hold onto its electrons, and is influenced by factors such as atomic size and the effective nuclear charge experienced by the electrons.

First ionization energy generally increases across a period from left to right due to a stronger attraction between the electrons and the nucleus as the nuclear charge increases. Additionally, first ionization energy decreases down a group because the outermost electrons are further from the nucleus, offset by the additional electron shells, which cause effective nuclear charge to be felt less by the valence electrons.

This property impacts how atoms undergo chemical reactions, particularly those that involve the loss of an electron. Understanding the first ionization energy helps in predicting the reactivity of elements, which is essential for students studying chemical bonding and periodicity.