Complete The Following Table For The Designated Atoms

Completing a table with specific information for designated atoms is a fundamental exercise in chemistry, helping us understand their properties and behavior. This process often involves determining the number of protons, neutrons, and electrons, understanding electron configuration, identifying valence electrons, and potentially predicting the atom's common ions and bonding behavior. Let's delve into how to comprehensively complete such a table, using examples to illustrate the process.

Understanding the Atomic Table

Before filling the table, let's define the key terms:

• Atomic Number (Z): The number of protons in the nucleus of an atom. This defines the element.
• Mass Number (A): The total number of protons and neutrons in the nucleus of an atom.
• Number of Protons: Equal to the atomic number (Z).
• Number of Neutrons: Calculated by subtracting the atomic number (Z) from the mass number (A): A - Z.
• Number of Electrons: In a neutral atom, the number of electrons is equal to the number of protons (Z). If the atom is an ion (has a charge), the number of electrons will differ.
• Electron Configuration: The arrangement of electrons in the different energy levels and sublevels (orbitals) within an atom.
• Valence Electrons: The electrons in the outermost energy level of an atom. These are the electrons involved in chemical bonding.
• Common Ions: Atoms can gain or lose electrons to achieve a stable electron configuration. The resulting charged particle is called an ion.

Steps to Complete the Table

Here's a systematic approach to completing the table for designated atoms:

1. Identify the Element: Determine the element symbol from the given atom or its name.

2. Find the Atomic Number (Z): Use the periodic table to find the atomic number of the element.

3. Determine the Mass Number (A):

   • If the problem gives you a specific isotope (e.g., Carbon-14), the mass number is the number following the element name (14 in the example).
   • If only the element is given, you'll typically use the average atomic mass from the periodic table. Round this value to the nearest whole number to obtain the most common mass number. Note this isn't always appropriate, and the question may intend you to know the most stable/common isotope.

4. Calculate Protons, Neutrons, and Electrons:

   • Protons: Number of protons = Atomic Number (Z)
   • Neutrons: Number of neutrons = Mass Number (A) - Atomic Number (Z)
   • Electrons:
     • For a neutral atom: Number of electrons = Atomic Number (Z)
     • For an ion:
       • If the ion has a positive charge (cation), it has lost electrons. Subtract the charge from the atomic number.
       • If the ion has a negative charge (anion), it has gained electrons. Add the charge to the atomic number.

5. Write the Electron Configuration: Use the Aufbau principle, Hund's rule, and the Pauli exclusion principle to fill the electron shells and subshells in the correct order. A good way to remember the order is with the following diagram:

   1s
   2s 2p
   3s 3p 3d
   4s 4p 4d 4f
   5s 5p 5d 5f 5g
   6s 6p 6d 6f 6g 6h
   7s 7p 7d 7f 7g 7h 7i
   

   Remember the number of electrons each orbital can hold:

   • s orbital: 2 electrons
   • p orbital: 6 electrons
   • d orbital: 10 electrons
   • f orbital: 14 electrons

   Example: For Sodium (Na, Z=11), the electron configuration is 1s² 2s² 2p⁶ 3s¹.

6. Determine the Number of Valence Electrons: Identify the electrons in the outermost shell (highest energy level) of the electron configuration.

   • Example: For Sodium (Na), the outermost shell is the 3rd shell (3s¹), so it has 1 valence electron.

7. Predict Common Ions: Atoms gain or lose electrons to achieve a stable octet (8 valence electrons) or duet (2 valence electrons for elements like hydrogen and helium).

   • Metals tend to lose electrons to form positive ions (cations).
   • Nonmetals tend to gain electrons to form negative ions (anions).
   • Example: Sodium (Na) readily loses 1 electron to form Na⁺, achieving the electron configuration of Neon (Ne). Chlorine (Cl) readily gains 1 electron to form Cl⁻, achieving the electron configuration of Argon (Ar).

Examples and Worked Solutions

Let's work through some examples to illustrate the process. We'll use the following table format:

Example 1: Oxygen (O)

1. Element: Oxygen (O)
2. Atomic Number (Z): 8 (from the periodic table)
3. Mass Number (A): 16 (The most abundant isotope is Oxygen-16, which is typically implied if no isotope is specified. The average atomic mass from the periodic table is ~16.00)
4. Protons, Neutrons, and Electrons:
   • Protons: 8
   • Neutrons: 16 - 8 = 8
   • Electrons: 8 (neutral atom)
5. Electron Configuration: 1s² 2s² 2p⁴
6. Valence Electrons: 6 (2s² 2p⁴)
7. Common Ion: O²⁻ (Oxygen gains 2 electrons to achieve a stable octet, becoming isoelectronic with Neon)

Completed Table:

Example 2: Potassium Ion (K⁺)

1. Element: Potassium (K)
2. Atomic Number (Z): 19 (from the periodic table)
3. Mass Number (A): 39 (The most abundant isotope is Potassium-39, which is typically implied if no isotope is specified. The average atomic mass from the periodic table is ~39.10)
4. Protons, Neutrons, and Electrons:
   • Protons: 19
   • Neutrons: 39 - 19 = 20
   • Electrons: 18 (K⁺ has lost 1 electron because of the +1 charge)
5. Electron Configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ (Note: This is the same electron configuration as Argon)
6. Valence Electrons: 0 (The outermost shell is now the 3rd shell, which is full)
7. Common Ion: K⁺ (This is already the common ion as indicated in the problem)

Completed Table:

Example 3: Carbon-14 (¹⁴C)

1. Element: Carbon (C)
2. Atomic Number (Z): 6 (from the periodic table)
3. Mass Number (A): 14 (Given in the problem as Carbon-14)
4. Protons, Neutrons, and Electrons:
   • Protons: 6
   • Neutrons: 14 - 6 = 8
   • Electrons: 6 (neutral atom)
5. Electron Configuration: 1s² 2s² 2p²
6. Valence Electrons: 4 (2s² 2p²)
7. Common Ion: While carbon can form C⁴⁺ or C⁴⁻ ions, it more commonly forms covalent bonds. For the purpose of this exercise, we'll consider the more common scenario of covalent bonding and indicate N/A.

Completed Table:

Example 4: Iron (Fe)

1. Element: Iron (Fe)
2. Atomic Number (Z): 26 (from the periodic table)
3. Mass Number (A): 56 (The most abundant isotope is Iron-56, which is typically implied if no isotope is specified. The average atomic mass from the periodic table is ~55.85)
4. Protons, Neutrons, and Electrons:
   • Protons: 26
   • Neutrons: 56 - 26 = 30
   • Electrons: 26 (neutral atom)
5. Electron Configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶
6. Valence Electrons: 8 (2 from the 4s orbital and 6 from the 3d orbitals) Note: It's important to recognize that for transition metals, both the s and d electrons in the outermost shells are considered valence electrons.
7. Common Ion: Fe²⁺, Fe³⁺ (Iron can lose 2 or 3 electrons)

Completed Table:

Common Mistakes to Avoid

• Confusing Atomic Number and Mass Number: Remember that the atomic number defines the element, while the mass number is the sum of protons and neutrons.
• Incorrectly Calculating Neutrons: Always subtract the atomic number from the mass number to find the number of neutrons.
• Forgetting to Adjust Electrons for Ions: Cations have fewer electrons than the neutral atom, while anions have more.
• Incorrect Electron Configuration: Pay close attention to the order of filling orbitals and Hund's rule.
• Misidentifying Valence Electrons: Remember to consider all electrons in the outermost shell, including s and p electrons, and be particularly careful with transition metals and d-orbital electrons.
• Assuming Average Atomic Mass is Always the Mass Number: While useful, if a specific isotope is given, use that isotope's mass number.

Importance of This Exercise

Completing these tables is not just about memorizing numbers. It helps you:

• Understand the Structure of Atoms: It reinforces your understanding of the subatomic particles that make up atoms.
• Predict Chemical Behavior: By knowing the number of valence electrons, you can predict how an atom will interact with other atoms to form chemical bonds.
• Connect Microscopic Properties to Macroscopic Behavior: The properties of individual atoms ultimately determine the properties of the matter we see and interact with every day.
• Foundation for Advanced Chemistry: This knowledge is crucial for understanding more advanced topics such as chemical bonding, reaction mechanisms, and spectroscopy.

Beyond the Basics: Advanced Considerations

While the above provides a solid foundation, here are some advanced points:

• Exceptions to Hund's Rule and Aufbau Principle: There are some exceptions to the filling order, particularly with transition metals like Chromium and Copper, where half-filled or fully-filled d-orbitals provide extra stability. For example, Chromium (Cr, Z=24) has an electron configuration of [Ar] 4s¹ 3d⁵, rather than the expected [Ar] 4s² 3d⁴.
• Isoelectronic Species: Ions and atoms that have the same electron configuration are called isoelectronic. For example, O²⁻, F⁻, Ne, Na⁺, and Mg²⁺ are all isoelectronic with the electron configuration 1s² 2s² 2p⁶.
• Relativistic Effects: For very heavy elements, relativistic effects (due to the high speeds of electrons near the nucleus) can significantly alter electron configurations and properties.
• Paramagnetism and Diamagnetism: Atoms or ions with unpaired electrons are paramagnetic (attracted to a magnetic field), while those with all paired electrons are diamagnetic (slightly repelled by a magnetic field). This can be predicted from the electron configuration.

Conclusion

Completing tables for designated atoms is a crucial exercise for solidifying your understanding of atomic structure and properties. By following the steps outlined above and practicing regularly, you can master this skill and build a strong foundation for further studies in chemistry. Remember to pay attention to details, avoid common mistakes, and always refer to the periodic table as your primary source of information. By meticulously working through these examples, you'll be well-equipped to tackle any similar problems and gain a deeper appreciation for the fundamental building blocks of matter.