Problem 51 - Chemistry: An Atoms-Focused Approach

Question

Draw Lewis symbols for \(\mathrm{K}^{+}, \mathrm{Al}^{3+}, \mathrm{N}^{3-},\) and \(\mathrm{I}^{-}.\)

Step-by-Step Solution

Answer

Question: Draw the Lewis symbols for the ions \(\mathrm{K}^{+}, \mathrm{Al}^{3+}, \mathrm{N}^{3-}, \mathrm{I}^{-}.\) Answer: The Lewis symbols for these ions are: 1. \(\mathrm{K}^{+}\): K^+ 2. \(\mathrm{Al}^{3+}\): Al^3+ 3. \(\mathrm{N}^{3-}\): :N:^: 4. \(\mathrm{I}^{-}\): :I::^:

Step 1: Find the number of valence electrons in each atom

To find the number of valence electrons in each atom, check their position in the periodic table. In the given exercise, we are looking at K, Al, N, and I. Their electron configurations and valence electrons are: 1. K (Potassium): [Ar] 4s1 - 1 Valence electron 2. Al (Aluminum): [Ne] 3s2 3p1 - 3 Valence electrons 3. N (Nitrogen): [He] 2s2 2p3 - 5 Valence electrons 4. I (Iodine): [Kr] 5s2 4d10 5p5 - 7 Valence electrons

Step 2: Determine the ion charge

As given in the exercise, our ions are \(\mathrm{K}^{+}, \mathrm{Al}^{3+}, \mathrm{N}^{3-}, \mathrm{I}^{-}.\) The charge of the ions shows that: 1. K atom has lost 1 electron to form \(\mathrm{K}^{+}\) ion 2. Al atom has lost 3 electrons to form \(\mathrm{Al}^{3+}\) ion 3. N atom has gained 3 electrons to form \(\mathrm{N}^{3-}\) ion 4. I atom has gained 1 electron to form \(\mathrm{I}^{-}\) ion

Step 3: Draw the Lewis symbols for the ions

We will now draw Lewis symbols for each ion: 1. \(\mathrm{K}^{+}\): K atom originally had 1 valence electron. It loses 1 electron to form \(\mathrm{K}^{+}\) ion, leaving it with no valence electrons. Thus, the Lewis symbol for \(\mathrm{K}^{+}\) is just K with a +1 charge: K^+ 2. \(\mathrm{Al}^{3+}\): Al atom originally had 3 valence electrons. It loses all 3 electrons to form \(\mathrm{Al}^{3+}\) ion, leaving it with no valence electrons. Thus, the Lewis symbol for \(\mathrm{Al}^{3+}\) is just Al with a +3 charge: Al^3+ 3. \(\mathrm{N}^{3-}\): N atom originally had 5 valence electrons. It gains 3 electrons to form \(\mathrm{N}^{3-}\) ion. Now, it has a total of 8 valence electrons. The Lewis symbol for \(\mathrm{N}^{3-}\) is N with 8 dots around it in pairs and a -3 charge: N^3- :N:^: 4. \(\mathrm{I}^{-}\): I atom originally had 7 valence electrons. It gains 1 electron to form \(\mathrm{I}^{-}\) ion. Now, it has a total of 8 valence electrons. The Lewis symbol for \(\mathrm{I}^{-}\) is I with 8 dots around it in pairs and a -1 charge: I^- :I::^:

Key Concepts

Valence Electrons

Valence electrons are crucial for understanding chemical bonding and reactions. They are the electrons located in the outermost shell of an atom. These electrons are responsible for interacting with other atoms. In chemistry, the behavior of an atom during bonding is primarily dictated by its valence electrons.
When identifying valence electrons, you can look at the group number in the periodic table for main-group elements. For example:

• Potassium (K) is in group 1 and has 1 valence electron.
• Aluminum (Al) is in group 13 and has 3 valence electrons.
• Nitrogen (N) is in group 15 and has 5 valence electrons.
• Iodine (I) is in group 17 and has 7 valence electrons.

Atoms strive to achieve a stable electron configuration, often resembling the nearest noble gas, by either gaining, losing, or sharing valence electrons. This pursuit of stability is the key driver of chemical bonding and the formation of ions.

Ions

Ions are charged particles formed when atoms gain or lose electrons. This electron transfer results in the atom acquiring a positive or negative charge. Atoms become ions to achieve a more stable electron configuration. Here are some examples of ion formation, relevant to our exercise:

• Cations: Positively charged ions formed when atoms lose electrons.
  - Potassium Ion ( K^{+} ): Formed when K loses 1 electron.
• Aluminum Ion ( Al^{3+} ): Formed when Al loses 3 electrons.


• Anions: Negatively charged ions formed when atoms gain electrons.
  - Nitride Ion ( N^{3-} ): Formed when N gains 3 electrons.
• Iodide Ion ( I^{-} ): Formed when I gains 1 electron.

Understanding ions is fundamental in predicting the behavior of substances during chemical reactions, as ions typically participate in ionic bonds.

Electron Configuration

Electron configuration describes the distribution of electrons in an atom's orbitals. It is a notation that marks the position of electrons around the nucleus of the atom. For students learning chemistry, understanding electron configurations helps in determining the chemical properties and reactivity of an element.
The electron configuration directly affects the number of valence electrons an atom has. Consider the following configurations for our specific atoms:

• Potassium (K): [Ar] 4s¹ – Meaning it has 1 valence electron in the 4s orbital.
• Aluminum (Al): [Ne] 3s² 3p¹ – 3 valence electrons in the 3rd shell orbitals.
• Nitrogen (N): [He] 2s² 2p³ – 5 valence electrons.
• Iodine (I): [Kr] 5s² 4d¹⁰ 5p⁵ – 7 valence electrons.

The configuration helps predict how atoms will achieve stable octets and participate in bonding by gaining, losing, or sharing electrons in chemical reactions.

Periodic Table

The periodic table is an essential tool for chemists. It organizes elements based on increasing atomic number and recurring chemical properties. Each element in the table can be identified with its electron configuration, valence electrons, and possible ion forms. The organization of the periodic table helps predict the behavior of different elements in reactions.
Elements are arranged in groups and periods:

• Groups: Vertical columns that indicate the number of valence electrons in main-group elements.
• Periods: Horizontal rows that show the number of electron shells.

For example, knowing that potassium is in group 1 informs you that it has 1 valence electron. Similarly, iodine in group 17 has 7 valence electrons. This structured form of information makes it easier to predict how elements will interact chemically to form compounds, including the creation of ions. By using the periodic table, students can intuitively grasp relationships among elements and foresee molecular formations and reactions.