Question

Write the electron configuration for each of the following atoms and for the simple ion that the element most commonly forms. In each case, indicate which noble gas has the same electron configuration as the ion. a. sodium, \(Z=11\) b. iodine, \(Z=53\) c. calcium, \(Z=20\) d. nitrogen, \(Z=7\) e. fluorine, \(Z=9\)

Short answer

a. Sodium: Electron configuration: \(1s^22s^22p^63s^1\); Na+ ion configuration: \(1s^22s^22p^6\); Noble gas: Neon (Ne) b. Iodine: Electron configuration: \(1s^22s^22p^63s^23p^63d^{10}4s^24p^64d^{10}5s^25p^3\); I- ion configuration: \(1s^22s^22p^63s^23p^63d^{10}4s^24p^64d^{10}5s^25p^4\); Noble gas: Xenon (Xe) c. Calcium: Electron configuration: \(1s^22s^22p^63s^23p^64s^2\); Ca2+ ion configuration: \(1s^22s^22p^63s^23p^6\); Noble gas: Argon (Ar) d. Nitrogen: Electron configuration: \(1s^22s^22p^3\); N3- ion configuration: \(1s^22s^22p^6\); Noble gas: Neon (Ne) e. Fluorine: Electron configuration: \(1s^22s^22p^5\); F- ion configuration: \(1s^22s^22p^6\); Noble gas: Neon (Ne)

Step-by-step solution

a. Sodium (Z=11) - Electron Configuration

We begin by finding the electron configuration for sodium, which has an atomic number (Z) of 11. Starting from hydrogen and filling the orbitals according to the periodic table, we have 1s, 2s, 2p, 3s, 3p, 4s, 3d... For sodium: 1. Fill the 1s orbital with 2 electrons \((1s^2)\) 2. Fill the 2s orbital with 2 electrons \((2s^2)\) 3. Fill the 2p orbital with 6 electrons \((2p^6)\) 4. 1 electron goes to the 3s orbital \((3s^1)\) Electron configuration: \(1s^22s^22p^63s^1\)

a. Sodium (Z=11) - Common Ion and Noble Gas

The most common ion for sodium is the positively charged ion Na+, which is formed when sodium loses 1 electron from the 3s orbital. This gives the sodium ion the electron configuration of \(1s^22s^22p^6\). This electron configuration is the same as that of the noble gas neon (Ne, Z=10).

b. Iodine (Z=53) - Electron Configuration

Now we will find the electron configuration for iodine with an atomic number (Z) of 53: 1. Fill the 1s orbital with 2 electrons \((1s^2)\) 2. Fill the 2s, 2p orbitals \((2s^22p^6)\) 3. Fill the 3s, 3p, 3d orbitals \((3s^23p^63d^{10})\) 4. Fill the 4s, 4p, 4d, 4f orbitals \((4s^24p^64d^{10})\) 5. Add 3 electrons to the 5s and 5p orbitals \((5s^25p^3)\) Electron configuration: \(1s^22s^22p^63s^23p^63d^{10}4s^24p^64d^{10}5s^25p^3\)

b. Iodine (Z=53) - Common Ion and Noble Gas

The most common ion for iodine is the negatively charged ion I-, which is formed when iodine gains 1 electron in the 5p orbital. This gives the iodine ion the electron configuration of \(1s^22s^22p^63s^23p^63d^{10}4s^24p^64d^{10}5s^25p^4\). This electron configuration is the same as that of the noble gas xenon (Xe, Z=54).

c. Calcium (Z=20) - Electron Configuration

Now for calcium with the atomic number (Z) of 20: 1. Fill the 1s, 2s, 2p orbitals \((1s^22s^22p^6)\) 2. Fill the 3s, 3p, 3d orbitals \((3s^23p^63d^0)\) 3. Add 2 electrons to the 4s orbital \((4s^2)\) Electron configuration: \(1s^22s^22p^63s^23p^64s^2\)

c. Calcium (Z=20) - Common Ion and Noble Gas

The most common ion for calcium is the positively charged ion Ca2+, which is formed when calcium loses 2 electrons from the 4s orbital. This gives the calcium ion the electron configuration of \(1s^22s^22p^63s^23p^6\). This electron configuration is the same as that of the noble gas argon (Ar, Z=18).

d. Nitrogen (Z=7) - Electron Configuration

For nitrogen with the atomic number (Z) of 7: 1. Fill the 1s, 2s orbitals (1s2 2s2) 2. Add 3 electrons to the 2p orbital (2p3) Electron configuration: \(1s^22s^22p^3\)

d. Nitrogen (Z=7) - Common Ion and Noble Gas

The most common ion for nitrogen is the negatively charged ion N3-, which is formed when nitrogen gains 3 electrons in the 2p orbital. This gives the nitrogen ion the electron configuration of \(1s^22s^22p^6\). This electron configuration is the same as that of the noble gas neon (Ne, Z=10).

e. Fluorine (Z=9) - Electron Configuration

Lastly, for fluorine with the atomic number (Z) of 9: 1. Fill the 1s, 2s orbitals \(1s^22s^2\) 2. Add 5 electrons to the 2p orbital \(2p^5\) Electron configuration: \(1s^22s^22p^5\)

e. Fluorine (Z=9) - Common Ion and Noble Gas

The most common ion for fluorine is the negatively charged ion F-, which is formed when fluorine gains 1 electron in the 2p orbital. This gives the fluorine ion the electron configuration of \(1s^22s^22p^6\). This electron configuration is the same as that of the noble gas neon (Ne, Z=10).

Key concepts

Atomic Number

The atomic number, represented by the symbol 'Z', is a fundamental property that defines each element in the periodic table. It refers to the number of protons in the nucleus of an atom, and by extension, also determines the number of electrons orbiting the nucleus in a neutral atom.

Understanding the atomic number is crucial when determining an element's electron configuration, as it tells us exactly how many electrons we need to distribute across the available orbitals. For example, sodium has an atomic number of 11, which means there are 11 electrons that need to be arranged in orbitals.

It is the atomic number that gives an element its unique identity and properties. As we move across the periodic table from left to right, the atomic number increases, signifying the addition of one proton and one electron for each subsequent element.

Noble Gas Configuration

Noble gas configuration is a highly stable electron arrangement that non-noble gas elements achieve through ion formation. Atoms tend to gain, lose, or share electrons to complete their outer shell and resemble the electron configuration of a noble gas, the inert elements located in Group 18 of the periodic table.

As demonstrated in the exercise, ions are formed when elements achieve a noble gas electron configuration. For example, when sodium (Na) loses an electron, it forms Na⁺ with the same electron configuration as neon (Ne). Similarly, by gaining an electron, fluorine (F) forms F^- and mimics the electron configuration of neon (Ne).

This concept is important in understanding chemical reactivity and bonding, as atoms are more stable and less reactive when they have achieved a noble gas configuration.

Electron Orbital Filling

Electron orbital filling is the process by which electrons occupy the orbitals around the nucleus of an atom. Electrons fill orbitals following specific rules: the Aufbau principle, the Pauli exclusion principle, and Hund's rule.

The Aufbau principle guides the filling order, from the lowest to the highest energy levels. According to this principle, the 1s orbital is filled first, followed by the 2s, 2p, and so on. This is why we see the electron configuration for sodium starting with 1s² before moving to higher energy orbitals.

The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers. Essentially, this means that an orbital can hold a maximum of two electrons with opposite spins. Hund's rule tells us that electrons will fill degenerate orbitals (orbitals with the same energy) singly before pairing up.

Electron orbital filling allows us to visualize how electrons are structured around an atom, influencing the atom's chemical properties and the types of bonds it can form.