Master Electron Configuration
Comprehensive study guide for understanding atomic structure and electron behavior
Overview
Master the fundamental principles of electron arrangement in atoms
- Aufbau principle: Electrons occupy the lowest-energy available subshells first.
- Hund's rule: In a set of degenerate orbitals (p, d, f), place one electron in each with the same spin before pairing.
- Pauli exclusion principle: Each orbital holds at most two electrons and they must have opposite spins (↑↓).
- Exceptions: Special cases like Chromium (Cr) and Copper (Cu)
- Ion Formation: How atoms gain or lose electrons to form ions
- Octet Rule: Why atoms seek 8 valence electrons for stability
Core Rules
The three fundamental principles that govern electron behavior
- Aufbau principle: Electrons occupy the lowest-energy available subshells first.
- Hund's rule: In a set of degenerate orbitals (p, d, f), place one electron in each with the same spin before pairing.
- Pauli exclusion principle: Each orbital holds at most two electrons and they must have opposite spins (↑↓).
- Energy Order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p...
- Diagonal Pattern: Follow the energy level progression
Exceptions
Special cases where elements break the standard rules for stability
- Chromium (Cr): [Ar] 4s¹ 3d⁵ instead of [Ar] 4s² 3d⁴
- Copper (Cu): [Ar] 4s¹ 3d¹⁰ instead of [Ar] 4s² 3d⁹
- Molybdenum (Mo): [Kr] 5s¹ 4d⁵ (half-filled d-subshell)
- Silver (Ag): [Kr] 5s¹ 4d¹⁰ (filled d-subshell)
- Why Exceptions: Half-filled and filled subshells provide extra stability
Ion Formation
Understanding how atoms gain and lose electrons to form ions
- Cations: Positively charged ions formed by losing electrons
- Anions: Negatively charged ions formed by gaining electrons
- Main Group Metals: Lose electrons from outermost shell
- Transition Metals: Lose s electrons first, then d electrons
- Non-metals: Gain electrons to achieve noble gas configuration
Practice
Interactive exercises to master electron configuration
- Interactive Atom Builder: Build atoms step by step
- Guided Problem Solving: Step-by-step guidance
- Interactive Quizzes: Test your knowledge
- Simulation Activities: Visualize electron behavior
- Practice Problems: Master the concepts
Assessment
Electron configuration focused examinations
- Paper 1: Core Principles & Rules (50 marks)
- Paper 2: Exceptions & Complex Cases (50 marks)
- Paper 3: Ion Formation & Applications (60 marks)
- Mark Schemes: Detailed Edexcel-style marking criteria
- Timed Exams: Focused electron configuration assessments
Overview - Master Electron Configuration
Welcome to your comprehensive journey into the world of electron configurations. This overview will introduce you to the fundamental concepts that govern how electrons arrange themselves in atoms.
Core Concepts Overview
Essential principles that govern how electrons arrange themselves in atoms. These are the exact same definitions used in the main app's study and game modes.
Aufbau principle
Electrons occupy the lowest-energy available subshells first.
Hund's rule
In a set of degenerate orbitals (p, d, f), place one electron in each with the same spin before pairing.
Pauli exclusion principle
Each orbital holds at most two electrons and they must have opposite spins (↑↓).
Exceptions
Special cases like Chromium and Copper
Ion Formation
How atoms gain or lose electrons
Octet Rule
Atoms seek 8 valence electrons for stability
Your Learning Journey
Start your journey through electron configuration concepts and track your progress!
Learning Path
Follow the structured learning path above. Click on concept tiles to learn more and track your progress. Each step builds upon the previous one, ensuring a solid foundation in electron configuration.
Your Learning Progress
Track your progress through the electron configuration concepts and unlock achievements as you learn!
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Electron Configuration Specifications
Key learning objectives for mastering electron configuration principles
📚 Core Principles
- Write electron configurations for elements up to Z=36
- Understand the Aufbau principle and energy level order
- Apply Hund's rule to degenerate orbitals
- Explain the Pauli exclusion principle and opposite spins
🎯 Advanced Concepts
- Explain exceptions like Cr and Cu configurations
- Understand ion formation and oxidation states
- Apply the octet rule to predict stability
- Relate configuration to periodic trends
🔬 Practical Skills
- Use noble gas notation for configurations
- Predict chemical properties from configurations
Key Vocabulary
Essential terms and definitions for electron configuration
Why Electron Configurations Matter
Electron configurations determine an element's chemical properties, reactivity, and bonding behavior. Understanding these principles is essential for:
Chemical Reactions
Predict how elements will react with each other
Periodic Trends
Understand patterns across the periodic table
Bonding Patterns
Explain how atoms form molecules
Material Design
Create new materials with desired properties
Your Learning Progress
Ion Formation
Understanding how atoms gain and lose electrons to form ions
Learning Objectives
By the end of this section, you will be able to:
🎯 Cation Formation
Explain how metals lose electrons to form positive ions
- Understand electron loss from s and d orbitals
- Predict ion charges for main group metals
- Explain transition metal ion formation
⚡ Anion Formation
Describe how non-metals gain electrons to form negative ions
- Understand electron gain to achieve noble gas configuration
- Predict ion charges for non-metals
- Explain ionic radius changes
🔬 Transition Metals
Analyze the complex ion formation of transition metals
- Understand s-orbital electron loss first
- Explain variable oxidation states
- Predict common ion charges
Essential Vocabulary
An atom or molecule with a net electric charge due to the loss or gain of electrons
A positively charged ion formed when an atom loses electrons
A negatively charged ion formed when an atom gains electrons
The size of an ion, which changes when electrons are gained or lost
The charge an atom would have if all bonds were ionic
The stable electron configuration of noble gases (full outer shell)
Ions or atoms with the same number of electrons
Transition metals that can form ions with different charges
🔬 Cation Formation
Cations are formed when atoms lose electrons. This typically occurs with metals, which have relatively low ionization energies.
Main Group Metals
Lose electrons from their outermost shell to achieve noble gas configuration
Na: 1s² 2s² 2p⁶ 3s¹ → Na⁺: 1s² 2s² 2p⁶
Loses 1 electron from 3s orbital
Transition Metals
Lose electrons from s orbitals first, then d orbitals
Fe: [Ar] 4s² 3d⁶ → Fe²⁺: [Ar] 3d⁶
Loses 2 electrons from 4s orbital
⚡ Anion Formation
Anions are formed when atoms gain electrons. This typically occurs with non-metals, which have high electron affinities.
Halogens
Gain 1 electron to achieve noble gas configuration
Cl: [Ne] 3s² 3p⁵ → Cl⁻: [Ne] 3s² 3p⁶
Gains 1 electron in 3p orbital
Chalcogens
Gain 2 electrons to achieve noble gas configuration
O: [He] 2s² 2p⁴ → O²⁻: [He] 2s² 2p⁶
Gains 2 electrons in 2p orbital
🎮 Interactive Ion Formation Demonstrations
Explore how different elements form ions by gaining or losing electrons.
🔬 Cation Formation: Sodium (Na → Na⁺)
Electron Configuration
⚡ Anion Formation: Chlorine (Cl → Cl⁻)
Electron Configuration
🔬 Transition Metal: Iron (Fe → Fe²⁺)
Electron Configuration
📝 Vocabulary Quiz
Test your understanding of ion formation terminology.
Question 1: What is a cation?
Question 2: How do transition metals typically lose electrons?
Question 3: What happens to ionic radius when an atom becomes a cation?
Question 4: Which element forms a -2 anion?
Question 5: What is the electron configuration of Na⁺?
🧠 Quick Understanding Check
Apply your knowledge of ion formation to solve these problems.
Question 1: Predict the ion formed by magnesium (Mg)
Given: Mg has electron configuration [Ne] 3s²
Question 2: Which ion has the same electron configuration as Argon?
Question 3: What is the electron configuration of Fe³⁺?
Given: Fe has electron configuration [Ar] 4s² 3d⁶
⚠️ Electron Configuration Exceptions
Understanding special cases where elements break standard rules for enhanced stability
🎯 Learning Objectives
Exception Recognition
Identify which elements have electron configuration exceptions
Stability Principles
Understand why half-filled and filled subshells provide extra stability
Pattern Recognition
Recognize the patterns in d-block exceptions (Cr, Cu, Mo, Ag, etc.)
Configuration Writing
Write correct electron configurations for exception elements
📚 Essential Vocabulary
Exception Terms
Stability Terms
Configuration Terms
🔍 Exception Elements Overview
These elements break the standard Aufbau principle to achieve greater stability through half-filled or filled subshells:
Chromium (Cr) - Z=24
Why: Half-filled 3d subshell (5 electrons) provides extra stability
Copper (Cu) - Z=29
Why: Filled 3d subshell (10 electrons) provides maximum stability
Molybdenum (Mo) - Z=42
Why: Half-filled 4d subshell provides extra stability
Silver (Ag) - Z=47
Why: Filled 4d subshell provides maximum stability
🎮 Interactive Exception Demonstrations
Chromium (Cr) Exception
Watch how Chromium achieves stability by moving one electron from 4s to 3d:
Configuration
Copper (Cu) Exception
Watch how Copper achieves stability by moving one electron from 4s to 3d:
Configuration
🧠 Vocabulary Quiz
Question 1: What is an electron configuration exception?
Question 2: Why does Chromium have the configuration [Ar] 4s¹ 3d⁵?
Question 3: What configuration does Copper (Cu) actually have?
Question 4: What provides extra stability in exception elements?
⚡ Quick Understanding Check
Question 1: Which element has the exception configuration [Ar] 4s¹ 3d⁵?
Question 2: Why do exception elements break the Aufbau principle?
Question 3: What is the actual configuration of Silver (Ag)?
Question 4: How many electrons are in a half-filled d subshell?
Question 5: Which of these is NOT an exception element?
Quiz Results
Score: 0/5
A-Level Edexcel Pearson Assessment
Comprehensive examination-style assessments with detailed mark schemes
📋 Electron Configuration Assessment Structure
Comprehensive assessments focused specifically on electron configuration topics:
📄 Paper 1: Core Principles & Rules
Duration: 45 minutes | Marks: 50
- Aufbau Principle & Orbital Filling
- Hund's Rule & Electron Spin
- Pauli Exclusion Principle
- Orbital Diagrams & Notation
- Energy Level Ordering
🧪 Paper 2: Exceptions & Complex Cases
Duration: 45 minutes | Marks: 50
- Transition Metal Exceptions (Cr, Cu, Mo, Ag)
- Half-filled & Filled Subshell Stability
- d-block Configuration Patterns
- Exception Recognition & Explanation
- Stability Factors
🔬 Paper 3: Ion Formation & Applications
Duration: 60 minutes | Marks: 60
- Cation Formation (Main Group & Transition Metals)
- Anion Formation & Noble Gas Configuration
- Isoelectronic Species
- Ionic Radius Changes
- Magnetic Properties Prediction
📝 Electron Configuration Question Types
2-4 marks each - Write electron configurations for given elements, including exceptions.
3-6 marks each - Draw orbital diagrams showing electron arrangement and spin.
4-8 marks each - Explain why certain elements have unusual configurations.
3-6 marks each - Predict and write configurations for ions formed from neutral atoms.
5-10 marks each - Compare configurations, identify patterns, and analyze trends.
4-8 marks each - Apply configuration knowledge to predict properties and behavior.
🎯 Select Your Assessment
Choose a paper to begin your Edexcel-style examination practice:
📄 Paper 1: Core Principles & Rules
Master the fundamental principles of electron configuration
• Aufbau principle application
• Hund's rule & electron spin
• Pauli exclusion principle
• Orbital diagrams & notation
🧪 Paper 2: Exceptions & Complex Cases
Understand and explain electron configuration exceptions
• Transition metal exceptions (Cr, Cu)
• Half-filled & filled subshell stability
• d-block configuration patterns
• Exception recognition & explanation
🔬 Paper 3: Ion Formation & Applications
Apply electron configuration knowledge to ion formation
• Cation & anion formation
• Noble gas configuration
• Isoelectronic species
• Magnetic properties prediction
📚 Sample Questions Preview
Here are examples of the question types you'll encounter in each paper:
📄 Paper 1 Sample Question
Question 1 (a) [2 marks]
Write the electron configuration of chromium (Cr) and explain why it differs from the expected configuration based on the Aufbau principle.
M1: [Ar] 4s¹ 3d⁵ (1 mark)
M2: Half-filled 3d subshell provides extra stability (1 mark)
🧪 Paper 2 Sample Question
Question 2 (b) [4 marks]
Explain why copper (Cu) has the electron configuration [Ar] 4s¹ 3d¹⁰ instead of the expected [Ar] 4s² 3d⁹ configuration.
M1: Filled 3d subshell provides extra stability (1 mark)
M2: Energy difference between 4s and 3d orbitals is small (1 mark)
M3: Filled subshells have lower energy than partially filled ones (1 mark)
M4: This is an exception to the Aufbau principle (1 mark)
🔬 Paper 3 Sample Question
Question 3 (c) [5 marks]
Iron (Fe) has the electron configuration [Ar] 4s² 3d⁶. Write the electron configuration for Fe²⁺ and Fe³⁺ ions, explaining the order of electron loss.
M1: Fe²⁺: [Ar] 3d⁶ (1 mark)
M2: Fe³⁺: [Ar] 3d⁵ (1 mark)
M3: 4s electrons are lost first (1 mark)
M4: 4s orbitals are higher in energy than 3d orbitals (1 mark)
M5: Then 3d electrons are lost to form Fe³⁺ (1 mark)
📋 Assessment Instructions
Before You Begin:
- Time Management: Allocate time based on mark allocation (1 mark ≈ 1.5 minutes)
- Answer Format: Follow Edexcel conventions for showing working and explanations
- Marking: Use the provided mark schemes to self-assess your responses
- Practice: Complete papers under timed conditions for realistic exam experience
Marking Guidelines:
- Accuracy: Correct answers earn full marks
- Working: Show all steps in calculations for partial credit
- Explanations: Use precise scientific terminology
- Units: Always include appropriate units in numerical answers
Paper 1: Core Principles & Rules
Duration: 45 minutes | Total Marks: 50
📋 Instructions
- Answer ALL questions in this paper
- Write your answers in the spaces provided
- Show all working for calculation questions
- Use appropriate scientific terminology
- Total marks for this paper: 50
Question 1
(a) [2 marks]
Write the electron configuration for the following elements:
(i) Sodium (Na) - atomic number 11
(ii) Chlorine (Cl) - atomic number 17
M1: Na: 1s² 2s² 2p⁶ 3s¹ (1 mark)
M2: Cl: 1s² 2s² 2p⁶ 3s² 3p⁵ (1 mark)
(b) [3 marks]
Draw the orbital diagram for nitrogen (N) showing the arrangement of electrons in the 2p subshell. Use arrows to represent electrons and indicate their spin.
M1: Three separate 2p orbitals shown (1 mark)
M2: One electron in each orbital (1 mark)
M3: All electrons with same spin (parallel spins) (1 mark)
Accept: ↑ ↑ ↑ or ↓ ↓ ↓ in separate boxes
(c) [4 marks]
Explain the three fundamental principles that govern electron configuration:
(i) Aufbau principle
(ii) Hund's rule
(iii) Pauli exclusion principle
M1: Aufbau: Electrons occupy lowest energy orbitals first (1 mark)
M2: Hund's: In degenerate orbitals, electrons occupy separate orbitals with same spin before pairing (1 mark)
M3: Pauli: Maximum 2 electrons per orbital with opposite spins (1 mark)
M4: All three principles correctly explained (1 mark)
Question 2
(a) [3 marks]
Write the electron configuration for the following transition metals:
(i) Scandium (Sc) - atomic number 21
(ii) Titanium (Ti) - atomic number 22
(iii) Vanadium (V) - atomic number 23
M1: Sc: [Ar] 4s² 3d¹ (1 mark)
M2: Ti: [Ar] 4s² 3d² (1 mark)
M3: V: [Ar] 4s² 3d³ (1 mark)
(b) [4 marks]
Explain why the 4s orbital is filled before the 3d orbital in transition metals, despite 3d having a lower principal quantum number.
M1: 4s orbital has lower energy than 3d orbital (1 mark)
M2: Due to penetration effect of s orbitals (1 mark)
M3: s orbitals experience less shielding from inner electrons (1 mark)
M4: This follows the Aufbau principle (1 mark)
Question 3
(a) [5 marks]
Complete the following table by writing the electron configuration for each element:
| Element | Atomic Number | Electron Configuration |
|---|---|---|
| Magnesium (Mg) | 12 | |
| Aluminum (Al) | 13 | |
| Silicon (Si) | 14 | |
| Phosphorus (P) | 15 | |
| Sulfur (S) | 16 |
M1: Mg: 1s² 2s² 2p⁶ 3s² (1 mark)
M2: Al: 1s² 2s² 2p⁶ 3s² 3p¹ (1 mark)
M3: Si: 1s² 2s² 2p⁶ 3s² 3p² (1 mark)
M4: P: 1s² 2s² 2p⁶ 3s² 3p³ (1 mark)
M5: S: 1s² 2s² 2p⁶ 3s² 3p⁴ (1 mark)
(b) [3 marks]
Identify which element from the table above has the highest first ionization energy and explain your reasoning.
M1: Phosphorus (P) (1 mark)
M2: Has half-filled 3p subshell (1 mark)
M3: Half-filled subshell provides extra stability (1 mark)
Question 4
(a) [4 marks]
Draw the orbital diagram for oxygen (O) and explain how Hund's rule applies to the 2p subshell.
M1: Three 2p orbitals shown (1 mark)
M2: Two electrons in separate orbitals with same spin (1 mark)
M3: Hund's rule: electrons occupy separate orbitals first (1 mark)
M4: Electrons have same spin before pairing (1 mark)
(b) [3 marks]
Compare the electron configurations of oxygen (O) and fluorine (F). Explain why fluorine has a higher electronegativity.
M1: O: [He] 2s² 2p⁴, F: [He] 2s² 2p⁵ (1 mark)
M2: F has higher nuclear charge (1 mark)
M3: F is closer to noble gas configuration (1 mark)
Question 5
(a) [6 marks]
Write the electron configuration for the following ions and explain the process of ion formation:
(i) Na⁺ (from sodium)
(ii) Cl⁻ (from chlorine)
(iii) Mg²⁺ (from magnesium)
M1: Na⁺: 1s² 2s² 2p⁶ (1 mark)
M2: Cl⁻: 1s² 2s² 2p⁶ 3s² 3p⁶ (1 mark)
M3: Mg²⁺: 1s² 2s² 2p⁶ (1 mark)
M4: Na loses 1 electron from 3s orbital (1 mark)
M5: Cl gains 1 electron in 3p orbital (1 mark)
M6: Mg loses 2 electrons from 3s orbital (1 mark)
(b) [4 marks]
Explain why the ionic radius of Na⁺ is smaller than the atomic radius of Na, while the ionic radius of Cl⁻ is larger than the atomic radius of Cl.
M1: Na⁺ smaller: electron lost, less electron-electron repulsion (1 mark)
M2: Na⁺: same nuclear charge, fewer electrons (1 mark)
M3: Cl⁻ larger: electron gained, more electron-electron repulsion (1 mark)
M4: Cl⁻: same nuclear charge, more electrons (1 mark)
Paper 2: Exceptions & Complex Cases
Duration: 45 minutes | Total Marks: 50
📋 Instructions
- Answer ALL questions in this paper
- Focus on transition metal exceptions and complex cases
- Explain stability factors and energy considerations
- Use appropriate scientific terminology
- Total marks for this paper: 50
Question 1
(a) [4 marks]
Write the electron configuration for the following transition metal exceptions:
(i) Chromium (Cr) - atomic number 24
(ii) Copper (Cu) - atomic number 29
(iii) Molybdenum (Mo) - atomic number 42
(iv) Silver (Ag) - atomic number 47
M1: Cr: [Ar] 4s¹ 3d⁵ (1 mark)
M2: Cu: [Ar] 4s¹ 3d¹⁰ (1 mark)
M3: Mo: [Kr] 5s¹ 4d⁵ (1 mark)
M4: Ag: [Kr] 5s¹ 4d¹⁰ (1 mark)
(b) [6 marks]
Explain why these elements have exceptional electron configurations instead of following the standard Aufbau principle. Include specific reasons for each element.
M1: Cr & Mo: Half-filled d subshell provides extra stability (1 mark)
M2: Cu & Ag: Filled d subshell provides extra stability (1 mark)
M3: Energy difference between s and d orbitals is small (1 mark)
M4: Symmetrical electron distribution reduces repulsion (1 mark)
M5: Lower energy configuration than expected Aufbau arrangement (1 mark)
M6: All four elements correctly explained (1 mark)
Question 2
(a) [5 marks]
Compare the electron configurations of the following pairs and explain the differences:
(i) Vanadium (V) vs Chromium (Cr)
(ii) Nickel (Ni) vs Copper (Cu)
(iii) Palladium (Pd) vs Silver (Ag)
M1: V: [Ar] 4s² 3d³, Cr: [Ar] 4s¹ 3d⁵ - Cr has half-filled d subshell (1 mark)
M2: Ni: [Ar] 4s² 3d⁸, Cu: [Ar] 4s¹ 3d¹⁰ - Cu has filled d subshell (1 mark)
M3: Pd: [Kr] 5s² 4d⁸, Ag: [Kr] 5s¹ 4d¹⁰ - Ag has filled d subshell (1 mark)
M4: Exceptions occur when stability gained exceeds energy cost (1 mark)
M5: All three comparisons correctly explained (1 mark)
(b) [4 marks]
Predict the electron configuration of Gold (Au) - atomic number 79 - and explain your reasoning.
M1: Au: [Xe] 6s¹ 5d¹⁰ (1 mark)
M2: Follows same pattern as Cu and Ag (1 mark)
M3: Filled 5d subshell provides extra stability (1 mark)
M4: Energy difference between 6s and 5d is small (1 mark)
Question 3
(a) [6 marks]
Complete the following table for transition metal ions, showing how they form from their neutral atoms:
| Element | Neutral Configuration | Ion Formed | Ion Configuration |
|---|---|---|---|
| Chromium (Cr) | [Ar] 4s¹ 3d⁵ | ||
| Copper (Cu) | [Ar] 4s¹ 3d¹⁰ | ||
| Iron (Fe) | [Ar] 4s² 3d⁶ |
M1: Cr³⁺: [Ar] 3d³ (1 mark)
M2: Cu²⁺: [Ar] 3d⁹ (1 mark)
M3: Fe²⁺: [Ar] 3d⁶ (1 mark)
M4: 4s electrons lost first (1 mark)
M5: Then 3d electrons lost (1 mark)
M6: All three configurations correct (1 mark)
(b) [3 marks]
Explain why transition metals lose 4s electrons before 3d electrons when forming ions.
M1: 4s electrons are higher in energy than 3d electrons (1 mark)
M2: 4s electrons are further from nucleus (1 mark)
M3: Less energy required to remove 4s electrons (1 mark)
Question 4
(a) [5 marks]
Identify which of the following elements are exceptions to the Aufbau principle and explain why:
Elements: Scandium (Sc), Titanium (Ti), Vanadium (V), Chromium (Cr), Manganese (Mn), Iron (Fe), Cobalt (Co), Nickel (Ni), Copper (Cu), Zinc (Zn)
M1: Exceptions: Cr and Cu (1 mark)
M2: Cr: half-filled 3d subshell provides stability (1 mark)
M3: Cu: filled 3d subshell provides stability (1 mark)
M4: Other elements follow standard Aufbau principle (1 mark)
M5: Stability gained exceeds energy cost of exception (1 mark)
(b) [4 marks]
Explain why elements like Scandium (Sc) and Zinc (Zn) do NOT have exceptional configurations, even though they are transition metals.
M1: Sc: only one d electron, no stability gain from half-filled subshell (1 mark)
M2: Zn: filled d subshell already achieved with standard configuration (1 mark)
M3: No energy advantage to changing s/d electron arrangement (1 mark)
M4: Standard Aufbau principle gives lowest energy configuration (1 mark)
Question 5
(a) [6 marks]
Draw orbital diagrams for the following elements showing their actual electron configurations:
(i) Chromium (Cr) - show 4s and 3d orbitals
(ii) Copper (Cu) - show 4s and 3d orbitals
(iii) Manganese (Mn) - show 4s and 3d orbitals
M1: Cr: 4s¹ (↑) 3d⁵ (↑ ↑ ↑ ↑ ↑) - one electron in 4s, five in separate 3d orbitals (2 marks)
M2: Cu: 4s¹ (↑) 3d¹⁰ (↑↓ ↑↓ ↑↓ ↑↓ ↑↓) - one electron in 4s, ten in 3d orbitals (2 marks)
M3: Mn: 4s² (↑↓) 3d⁵ (↑ ↑ ↑ ↑ ↑) - two electrons in 4s, five in separate 3d orbitals (2 marks)
(b) [4 marks]
Explain the magnetic properties of these three elements based on their electron configurations.
M1: Cr: paramagnetic - unpaired electrons in 3d orbitals (1 mark)
M2: Cu: diamagnetic - all electrons paired in 3d orbitals (1 mark)
M3: Mn: paramagnetic - unpaired electrons in 3d orbitals (1 mark)
M4: Unpaired electrons create magnetic moment (1 mark)
Paper 3: Ion Formation & Applications
Duration: 60 minutes | Total Marks: 60
📋 Instructions
- Answer ALL questions in this paper
- Focus on ion formation mechanisms and practical applications
- Include detailed explanations and calculations where required
- Use appropriate scientific terminology and units
- Total marks for this paper: 60
Question 1
(a) [6 marks]
Explain the process of ion formation for the following elements and write their ion configurations:
(i) Sodium (Na) - atomic number 11
(ii) Chlorine (Cl) - atomic number 17
(iii) Calcium (Ca) - atomic number 20
(iv) Oxygen (O) - atomic number 8
M1: Na: loses 1 electron from 3s orbital → Na⁺: [Ne] (1 mark)
M2: Cl: gains 1 electron in 3p orbital → Cl⁻: [Ar] (1 mark)
M3: Ca: loses 2 electrons from 4s orbital → Ca²⁺: [Ar] (1 mark)
M4: O: gains 2 electrons in 2p orbital → O²⁻: [Ne] (1 mark)
M5: All achieve noble gas configuration (1 mark)
M6: Correct electron loss/gain explanations (1 mark)
(b) [4 marks]
Calculate the ionic radius ratio for the following ion pairs and explain the trend:
Ion Pairs: Na⁺/Cl⁻, Ca²⁺/O²⁻, K⁺/Br⁻
M1: Na⁺/Cl⁻ ≈ 0.52, Ca²⁺/O²⁻ ≈ 0.74, K⁺/Br⁻ ≈ 0.68 (2 marks)
M2: Cations smaller than anions due to nuclear charge (1 mark)
M3: Higher charge = smaller ionic radius (1 mark)
Question 2
(a) [8 marks]
Complete the following table for transition metal ion formation:
| Element | Neutral Config | Ion Formed | Ion Config | Magnetic Property |
|---|---|---|---|---|
| Iron (Fe) | [Ar] 4s² 3d⁶ | |||
| Iron (Fe) | [Ar] 4s² 3d⁶ | |||
| Copper (Cu) | [Ar] 4s¹ 3d¹⁰ | |||
| Chromium (Cr) | [Ar] 4s¹ 3d⁵ |
M1: Fe²⁺: [Ar] 3d⁶, paramagnetic (2 marks)
M2: Fe³⁺: [Ar] 3d⁵, paramagnetic (2 marks)
M3: Cu²⁺: [Ar] 3d⁹, paramagnetic (2 marks)
M4: Cr³⁺: [Ar] 3d³, paramagnetic (2 marks)
(b) [4 marks]
Explain why Fe³⁺ is more stable than Fe²⁺ in aqueous solution, considering electron configurations and charge density.
M1: Fe³⁺ has half-filled 3d subshell (1 mark)
M2: Half-filled subshell provides extra stability (1 mark)
M3: Higher charge density in Fe³⁺ (1 mark)
M4: Better hydration in aqueous solution (1 mark)
Question 3
(a) [6 marks]
Predict the electron configuration and properties of the following ions:
(i) Mn²⁺ (manganese ion)
(ii) Co³⁺ (cobalt ion)
(iii) Ni²⁺ (nickel ion)
For each ion, state: electron configuration, magnetic property, and likely color in solution.
M1: Mn²⁺: [Ar] 3d⁵, paramagnetic, pale pink (2 marks)
M2: Co³⁺: [Ar] 3d⁶, paramagnetic, blue (2 marks)
M3: Ni²⁺: [Ar] 3d⁸, paramagnetic, green (2 marks)
(b) [4 marks]
Explain why transition metal ions are often colored, while main group metal ions (like Na⁺, Ca²⁺) are colorless.
M1: Transition metals have partially filled d orbitals (1 mark)
M2: d-d electronic transitions absorb visible light (1 mark)
M3: Main group ions have filled/empty outer shells (1 mark)
M4: No d-d transitions possible in main group ions (1 mark)
Question 4
(a) [5 marks]
Calculate the effective nuclear charge (Zeff) for the following ions and explain the trend:
Ions: Na⁺, Mg²⁺, Al³⁺, Si⁴⁺
Use the formula: Zeff = Z - S (where Z = atomic number, S = shielding constant)
M1: Na⁺: Zeff = 11 - 2 = 9 (1 mark)
M2: Mg²⁺: Zeff = 12 - 2 = 10 (1 mark)
M3: Al³⁺: Zeff = 13 - 2 = 11 (1 mark)
M4: Si⁴⁺: Zeff = 14 - 2 = 12 (1 mark)
M5: Zeff increases with increasing charge (1 mark)
(b) [3 marks]
Explain how effective nuclear charge affects ionic radius and ionization energy.
M1: Higher Zeff = smaller ionic radius (1 mark)
M2: Higher Zeff = higher ionization energy (1 mark)
M3: Electrons held more tightly by nucleus (1 mark)
Question 5
(a) [6 marks]
Explain the formation of complex ions and their electron configurations:
(i) [Fe(CN)₆]³⁻ - hexacyanoferrate(III) ion
(ii) [Cu(NH₃)₄]²⁺ - tetraamminecopper(II) ion
(iii) [Co(H₂O)₆]²⁺ - hexaaquacobalt(II) ion
M1: [Fe(CN)₆]³⁻: Fe³⁺ + 6CN⁻, d²sp³ hybridization (2 marks)
M2: [Cu(NH₃)₄]²⁺: Cu²⁺ + 4NH₃, dsp² hybridization (2 marks)
M3: [Co(H₂O)₆]²⁺: Co²⁺ + 6H₂O, sp³d² hybridization (2 marks)
(b) [4 marks]
Predict the magnetic properties and colors of these complex ions, explaining your reasoning.
M1: [Fe(CN)₆]³⁻: paramagnetic, red (1 mark)
M2: [Cu(NH₃)₄]²⁺: paramagnetic, blue (1 mark)
M3: [Co(H₂O)₆]²⁺: paramagnetic, pink (1 mark)
M4: All have unpaired electrons in d orbitals (1 mark)
Question 6
(a) [8 marks]
Design an experiment to determine the electron configuration of an unknown transition metal ion in solution. Include:
(i) Experimental method
(ii) Measurements to be taken
(iii) How to interpret results
(iv) Safety considerations
M1: UV-Vis spectroscopy to measure d-d transitions (2 marks)
M2: Magnetic susceptibility measurements (2 marks)
M3: Compare with known standards (2 marks)
M4: Safety: gloves, goggles, fume hood (2 marks)
(b) [4 marks]
Explain how the results of your experiment would help identify the electron configuration of the unknown ion.
M1: UV-Vis peaks indicate d-d transition energies (1 mark)
M2: Magnetic data shows number of unpaired electrons (1 mark)
M3: Color indicates specific d orbital splitting (1 mark)
M4: Combined data gives complete electron configuration (1 mark)
Ready to Practice?
Put your knowledge to the test with interactive electron configuration challenges. Build atoms step by step and master the principles through hands-on experience.
Core Rules - Interactive Learning
Master the three fundamental principles through interactive demonstrations. These are the exact same rules used in the main app's study and game modes.
Learning Objectives
What you will learn and master in this Core Rules section
🎯 Aufbau Principle
- Understand that electrons fill lowest energy orbitals first
- Learn the correct filling order: 1s → 2s → 2p → 3s → 3p → 4s → 3d...
- Recognize energy levels from n=1 (bottom) to n=7 (top)
- Apply the principle to predict electron configurations
🎯 Hund's Rule
- Understand that electrons spread out in degenerate orbitals first
- Learn that all unpaired electrons have the same spin (↑)
- Recognize that pairing only occurs after all orbitals have one electron
- Apply the rule to p, d, and f subshells
🎯 Pauli Exclusion Principle
- Understand that each orbital can hold maximum 2 electrons
- Learn that paired electrons must have opposite spins (↑↓)
- Recognize the difference between filled and paired orbitals
- Apply the principle to prevent orbital overcrowding
🎯 Interactive Application
- Use interactive demonstrations to visualize each rule
- Practice applying all three rules together
- Test understanding through quizzes and assessments
- Master the vocabulary and terminology
Essential Vocabulary
Key terms used in the Core Rules demonstrations and quizzes
🎯 Core Rules
🌌 Orbital Structure
⚡ Electron Properties
🎮 Interactive Terms
Official Rule Definitions
These are the exact definitions used in the main app's study and game modes:
1. Aufbau principle
Electrons occupy the lowest-energy available subshells first.
2. Hund's rule
In a set of degenerate orbitals (p, d, f), place one electron in each with the same spin before pairing.
3. Pauli exclusion principle
Each orbital holds at most two electrons and they must have opposite spins (↑↓).
Interactive Aufbau Sequence
Watch electrons fill orbitals in order of increasing energy. Click the buttons to see the sequence in action!
Hund's Rule Demonstration
See how electrons distribute themselves in degenerate orbitals. Click the orbitals to understand parallel vs. paired spins!
Pauli Exclusion Principle
Test the rule that no two electrons can have identical quantum numbers. Try to place electrons and see what happens!
Vocabulary Quiz
Test your understanding of the key chemistry terms before proceeding to the interactive concepts
1. What does Hund's Rule state?
2. What are degenerate orbitals?
3. What is the Pauli Exclusion Principle?
4. What are unpaired electrons?
5. What is a lone pair?
6. What does the Aufbau Principle state?
7. How many orbitals are in a p subshell?
8. What is the ground state?
Quick Understanding Check
Test your understanding of the core rules with these interactive questions!
Question 1: What is the correct order for filling orbitals?
Question 2: How should electrons be placed in degenerate orbitals?
Question 3: What is the maximum number of electrons per orbital?
Question 4: In a 2p subshell, how should 4 electrons be arranged?
Question 5: What happens when electrons pair in an orbital?
Question 6: Which element has the electron configuration 1s² 2s² 2p⁶?
Question 7: How many unpaired electrons does oxygen (O) have in its ground state?
Question 8: What is the correct electron configuration for carbon (C)?
Question 9: Which rule explains why electrons spread out before pairing?
Question 10: How many orbitals are in a d subshell?
Question 11: How are the 2p electrons arranged in nitrogen (N)?
Quiz Results
Score: 0/11
Interactive Practice Center
Master electron configuration through hands-on interactive activities. Build atoms, solve problems, and test your understanding with engaging exercises.
Interactive Atom Builder
Build atoms step by step and see how electrons fill orbitals according to the Aufbau principle, Hund's rule, and Pauli exclusion principle.
Select an element to begin
Choose from the dropdown above
Electron Configuration:
Orbital Diagram:
Guided Problem Solving
Work through electron configuration problems with step-by-step guidance and hints.
Problem Title
Problem description will appear here
Interactive Quiz
Test your knowledge with interactive multiple-choice and fill-in-the-blank questions.
Electron Configuration Quiz
Test your understanding
Question will appear here
Quiz Results
Interactive Simulations
Visualize electron behavior and orbital filling with interactive simulations.
Practice Problems
Work through a variety of electron configuration problems to master the concepts.
Basic Problems
Simple electron configurations
Intermediate Problems
Transition metals and exceptions
Advanced Problems
Complex ions and applications