Atomic Structure Worksheet 2 Answer Key

Atomic Structure Worksheet 2: Answers and Explanations

This comprehensive guide provides answers and detailed explanations for a hypothetical "Atomic Structure Worksheet 2." Since I don't have access to a specific worksheet with that name, I've created a representative worksheet covering key concepts in atomic structure. This guide will cover various aspects, including atomic number, mass number, isotopes, ions, electron configuration, and orbital diagrams. We'll tackle each problem type systematically, enhancing your understanding of atomic structure.

Section 1: Basic Atomic Structure

This section focuses on fundamental atomic concepts like atomic number, mass number, protons, neutrons, and electrons.

Problem 1: Determine the number of protons, neutrons, and electrons in an atom of Carbon-12.

Answer:

• Atomic Number: Carbon's atomic number is 6 (found on the periodic table). This means it has 6 protons.
• Protons: 6
• Electrons: In a neutral atom, the number of electrons equals the number of protons. Therefore, it has 6 electrons.
• Mass Number: The mass number is given as 12. Mass number = protons + neutrons.
• Neutrons: 12 (mass number) - 6 (protons) = 6 neutrons.

Problem 2: An atom has 17 protons, 18 neutrons, and 17 electrons. Identify the element and write its isotopic notation.

Answer:

• Element Identification: The number of protons defines the element. An element with 17 protons is Chlorine (Cl).
• Isotopic Notation: The isotopic notation is written as ³⁵Cl₁₇. The superscript (35) is the mass number (protons + neutrons = 17 + 18 = 35), and the subscript (17) is the atomic number.

Problem 3: Explain the difference between atomic number and mass number.

Answer:

The atomic number represents the number of protons in an atom's nucleus. This number uniquely identifies the element. The mass number represents the total number of protons and neutrons in an atom's nucleus. It's the total mass of the nucleus, expressed in atomic mass units (amu). The mass number can vary for different isotopes of the same element.

Section 2: Isotopes and Ions

This section delves into isotopes (atoms of the same element with different numbers of neutrons) and ions (charged atoms).

Problem 4: Explain what isotopes are and give an example.

Answer: Isotopes are atoms of the same element that have the same number of protons but a different number of neutrons. This results in different mass numbers. For example, Carbon-12 (¹²C) and Carbon-14 (¹⁴C) are isotopes of carbon. Both have 6 protons, but Carbon-12 has 6 neutrons, while Carbon-14 has 8 neutrons.

Problem 5: What is an ion? How does a cation differ from an anion?

Answer: An ion is an atom or molecule that has gained or lost one or more electrons, giving it a net positive or negative charge. A cation is a positively charged ion (it has lost electrons), while an anion is a negatively charged ion (it has gained electrons).

Problem 6: What is the charge of an ion with 11 protons and 10 electrons? Identify the element.

Answer:

• Charge: The ion has 11 protons (+11 charge) and 10 electrons (-10 charge). The net charge is +1.
• Element: An element with 11 protons is Sodium (Na). Therefore, this is a sodium ion, Na⁺.

Section 3: Electron Configuration and Orbital Diagrams

This section focuses on the arrangement of electrons within an atom, using electron configurations and orbital diagrams.

Problem 7: Write the electron configuration for Oxygen (O).

Answer: Oxygen has 8 electrons. Its electron configuration is 1s²2s²2p⁴. This indicates that the first energy level (n=1) has 2 electrons in the 1s orbital, and the second energy level (n=2) has 2 electrons in the 2s orbital and 4 electrons in the 2p orbitals.

Problem 8: Draw the orbital diagram for Nitrogen (N).

Answer: Nitrogen has 7 electrons. The orbital diagram would show:

1s: ↑↓ 2s: ↑↓ 2p: ↑ ↑ ↑ (One electron in each 2p orbital before pairing)

Problem 9: Explain Hund's rule and the Pauli exclusion principle.

Answer:

• Hund's Rule: Hund's rule states that electrons will individually occupy each orbital within a subshell before doubling up in any one orbital. This minimizes electron-electron repulsion.
• Pauli Exclusion Principle: The Pauli exclusion principle states that a maximum of two electrons can occupy a single atomic orbital, and these two electrons must have opposite spins (one spin up, one spin down).

Problem 10: What is the difference between an orbital and a subshell?

Answer: A subshell is a group of atomic orbitals that have the same principal quantum number (n) and the same azimuthal quantum number (l). For example, the 2p subshell contains three 2p orbitals (2p_x, 2p_y, 2p_z). An orbital is a region of space within an atom where there is a high probability of finding an electron. Each orbital can hold a maximum of two electrons.

Section 4: More Complex Atomic Structures and Advanced Concepts

This section explores more challenging aspects of atomic structure, requiring a deeper understanding of quantum numbers and electron configurations.

Problem 11: Write the electron configuration for Iron (Fe).

Answer: Iron (Fe) has 26 electrons. Its electron configuration is 1s²2s²2p⁶3s²3p⁶4s²3d⁶. Note the filling order of the 4s and 3d subshells.

Problem 12: Explain the Aufbau principle.

Answer: The Aufbau principle (German for "building-up principle") states that electrons first fill the lowest energy levels available before occupying higher energy levels. This principle guides the filling order of electrons in orbitals.

Problem 13: What are quantum numbers, and what information do they provide?

Answer: Quantum numbers are a set of numbers that describe the properties of atomic orbitals and the electrons within them. The four main quantum numbers are:

• Principal Quantum Number (n): Describes the energy level of the electron (n = 1, 2, 3...). Higher n values mean higher energy levels.
• Azimuthal Quantum Number (l): Describes the shape of the orbital (l = 0 to n-1). l = 0 represents an s orbital, l = 1 represents a p orbital, l = 2 represents a d orbital, and so on.
• Magnetic Quantum Number (m_l): Describes the orientation of the orbital in space (m_l = -l to +l). For example, the p subshell (l=1) has three orbitals (m_l = -1, 0, +1).
• Spin Quantum Number (m_s): Describes the intrinsic angular momentum of the electron (m_s = +1/2 or -1/2), representing "spin up" or "spin down."

Problem 14: Predict the number of valence electrons for Sulfur (S).

Answer: Sulfur (S) has an atomic number of 16. Its electron configuration is 1s²2s²2p⁶3s²3p⁴. Valence electrons are in the outermost energy level (n=3 in this case). Therefore, Sulfur has 6 valence electrons (2 in 3s and 4 in 3p).

Section 5: Applying Atomic Structure Concepts

This section applies the concepts learned to solve problems involving chemical bonding and reactivity.

Problem 15: Explain how the number of valence electrons influences the chemical reactivity of an element.

Answer: The number of valence electrons determines an element's chemical reactivity. Elements tend to react in ways that allow them to achieve a stable electron configuration, often by gaining, losing, or sharing electrons to achieve a full outermost electron shell (octet rule for many elements).

Problem 16: Why are noble gases generally unreactive?

Answer: Noble gases have a full outermost electron shell (except for Helium, which has a full first energy level). This stable electron configuration makes them very unreactive, as they have little tendency to gain, lose, or share electrons.

Problem 17: Predict the likely charge of an ion formed by Potassium (K).

Answer: Potassium (K) has one valence electron. To achieve a stable octet, it is more likely to lose this one electron to form a +1 ion (K⁺).

This expanded answer key provides detailed explanations for a wide range of atomic structure problems. Remember to consult your periodic table as a valuable resource throughout your studies. By understanding these fundamental concepts, you will build a strong foundation for further exploration of chemistry. Further practice with various problem types will solidify your understanding and improve your problem-solving skills. Remember to always refer back to the definitions and rules to ensure a firm grasp on the concepts.