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Magazine Chapter 8: Problem 111
Explain, in terms of their electron configurations, why \(\mathrm{Fe}^{2+}\) is more easily oxidized to \(\mathrm{Fe}^{3+}\) than \(\mathrm{Mn}^{2+}\) is to \(\mathrm{Mn}^{3+}\)
Short Answer
Expert verified
Fe2+ is more easily oxidized to Fe3+ than Mn2+ is to Mn3+ because Fe2+ can achieve a more stable half-filled d orbital configuration by losing one more electron while Mn2+, which already has a stable half-filled d orbital, would become less stable by losing an electron to become Mn3+.
Step by step solution
01
Understand Electron Configurations
Electron configurations describe the distribution of electrons in an atom's electron shells. They are determined by the Aufbau Principle which states that electrons always fill the lowest energy levels first, the Pauli Exclusion Principle which states that each orbital can hold a maximum of two electrons with opposite spins, and Hund's rule which states that electrons will occupy empty orbitals of equal energy before pairing up.
02
Identify the Electron Configurations of Fe2+, Fe3+, Mn2+ and Mn3+
The electron configuration for Fe2+ is [Ar] 3d6 and for Fe3+ it is [Ar] 3d5. For Mn2+, it is [Ar] 3d5 and for Mn3+ it is [Ar] 3d4.
03
Understand the Stability of Half-Filled and Fully Filled Orbitals
Half-filled and fully filled orbitals are particularly stable configurations. This stability comes from the symmetrical distribution of electrons and the maximum amount of unpaired electrons. In this context, Fe2+, which has four unpaired electrons, can lose one more electron to become Fe3+ with a half-filled (hence, stable) d orbitals while Mn2+ is already in a stable state with half-filled d orbitals so Mn3+ (which implies loss of electron from Mn2+) leads to a less stable configuration.
04
Compare the Ease of Oxidation
Since Fe2+ can become more stable by losing one more electron, it is more easily oxidized to Fe3+, compared to Mn2+ which is already in a stable state and would become less stable by losing an electron to become Mn3+.
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Electron Configurations
When exploring the intricacies of chemistry, electron configurations provide a fundamental blueprint for understanding how atoms interact during chemical reactions. These configurations map out the electrons surrounding a nucleus and show us the order in which atoms fill their available energy levels, or orbitals. Think of it like organizing a bookshelf, where the books must be arranged in specific slots based on their sizes.
Some key points about electron configurations include:
Some key points about electron configurations include:
- They're designated by a combination of quantum numbers and letters representing the orbital types (s, p, d, f).
- Configuration patterns follow a predictive order based on increasing energy.
- They influence an atom's magnetic properties, ability to bond with other atoms, and reactivity in chemical processes such as oxidation-reduction reactions.
Aufbau Principle
The Aufbau Principle is a guiding rule that dictates electron configurations. In simple terms, it's like the 'first come, first served' of electron arrangements, meaning electrons fill up the lowest energy orbitals available first before moving on to higher ones. This principle creates the foundation for an atom's structure.
Here鈥檚 why it is important:
Here鈥檚 why it is important:
- It helps us predict electron placement for both neutral atoms and ions.
- The energy of orbitals increases from 1s to 2s to 2p to 3s, and so on, following a specific sequence.
- Understanding this concept is vital when you are determining why certain ions are more or less stable after losing an electron.
Pauli Exclusion Principle
Central to the framework of quantum mechanics, the Pauli Exclusion Principle provides the 'no double occupancy' rule for electrons in the same orbital. It states that two electrons can share the same orbital only if they have opposite spins, like dancers moving in perfect unison but in opposite directions.
Some points about the Pauli Exclusion Principle include:
Some points about the Pauli Exclusion Principle include:
- It explains why electrons in the same orbital must have different spins, which is distinguished by spin quantum numbers: +1/2 or -1/2.
- This principle ensures that each electron in an atom has a unique set of quantum numbers.
- It's a key factor in determining the electron configurations of atoms and ions and, consequently, their chemical characteristics.
Hund's Rule
Hund's Rule can be thought of as a 'spread out before pairing up' strategy adopted by electrons. It explains that electrons will fill degenerate, or equal energy, orbitals singly as far as possible before starting to pair up. This behavior is akin to passengers preferring to sit in separate empty rows in a bus before sharing seats.
Key aspects of Hund's rule include:
Key aspects of Hund's rule include:
- It reflects the electrons' natural repulsion from one another due to their negative charge.
- By occupying separate orbitals, the electrons remain as unpaired as possible, leading to maximum total spin. This ultimately impacts an atom鈥檚 stability and reactivity.
- Applying Hund's rule is crucial when predicting how atoms will interact, particularly in processes such as the formation of Fe^{3+} from Fe^{2+}, where orbital occupancy plays a role.
Oxidation-Reduction Reactions
Diving into the collaborative world of electrons, oxidation-reduction reactions, or redox reactions, showcase the intimate dance of electron transfer between atoms. They are a type of chemical reaction where one substance loses electrons (oxidation) and another gains electrons (reduction).
Here's what makes redox reactions fascinating:
Here's what makes redox reactions fascinating:
- They're fundamental to processes like cellular respiration, photosynthesis, and even the rusting of iron.
- Oxidation states denote the charge an atom would have if all bonds were ionic, helping chemists keep track of electrons during reactions.
- An atom's propensity to lose or gain electrons is heavily influenced by its electron configuration and the relative stability that configuration provides.
