3.2.2
Ionisation & Electronegativity
Trends in Ionisation Energy
Trends in Ionisation Energy
Trends are seen in ionisation energy across periods and down groups as the strength of the attraction between the nucleus and the electron being removed changes.
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1st ionisation energy
1st ionisation energy
- The 1st ionisation energy is defined as the energy required to remove one mole of electrons from one mole of gaseous atoms in their ground state. An example is shown for sodium below:
- Na(g) → Na+(g) + e- ΔH = +496kJ mol-1
- This is an endothermic process.
Group trend
Group trend
- Ionisation energies decrease (become less endothermic) going down a group as the electron being removed experiences the same effective nuclear charge but is in a higher principal energy level, which is further away from the nucleus.
- The effective nuclear charge stays constant as although the nuclear charge increases so does the number of shielding electrons.
Periodic trend
Periodic trend
- Ionisation energies generally increase (become more endothermic) across a period as the effective nuclear charge increases. Since the valence electrons experience a stronger nuclear attraction, they are more difficult to remove.
Dips in ionisation energy
Dips in ionisation energy
- Dips in the trend of increasing ionisation energy across a period are observed between group 2 to group 3 and group 15 and group 16.
Group 2 to 3
Group 2 to 3
- The first ionisation of magnesium (group 2) corresponds to an electron being removed from a 3s orbital, whereas the first electron in aluminium (group 3) is removed from a 3p orbital.
- 3p is at higher energy than 3s and so it requires less energy to remove.
- Therefore the 1st ionisation energy of aluminium is lower than the 1st ionisation energy of magnesium.
Group 15 to 16
Group 15 to 16
- The first ionisation of sulphur (group 16) corresponds to an electron being removed from a doubly occupied p orbital, whereas the first electron in phosphorus (group 15) is removed from a singly occupied p orbital.
- There is repulsion between the electrons in the doubly occupied p orbital and so less energy is required to remove the first electron.
- Therefore the 1st ionisation energy of sulphur is less than the 1st ionisation energy of phosphorus.
Trends in Electron Affinity
Trends in Electron Affinity
Trends are seen in electron affinity across periods and down groups as the strength of the attraction between the nucleus and the incoming electron changes.
1st electron affinity
1st electron affinity
- The 1st electron affinity is the energy change that occurs when one mole of electrons is gained by one mole of gaseous atoms to form one mole of gaseous anions. An example of chlorine is shown below:
- Cl(g) + e- → Cl-(g) ΔH = -349kJ mol-1
- 1st electron affinity is typically an exothermic process due to the favourable interaction between the positive nucleus and incoming electron.
Group trend
Group trend
- Electron affinity decreases (becomes less exothermic) on descending a group as the incoming electron is further away from the nucleus as atomic radius increases.
Periodic trend
Periodic trend
- Electron affinity increases (becomes more exothermic) on going across a period as atoms at the end of the period have incomplete energy levels and higher effective nuclear charges.
Trends in Electronegativity
Trends in Electronegativity
Trends are seen in electronegativity across periods and down groups as the strength of the attraction of the bonding electron pair by the nucleus changes.
Electronegativity
Electronegativity
- Electronegativity is the ability of an atom to attract electrons in a covalent bond.
Group trend
Group trend
- Electronegativity decreases on going down a group as the bonding electrons become further away from the nucleus (atomic radius increases) and therefore the attraction of the bonding electrons is decreased.
Periodic trend
Periodic trend
- Electronegativity increases on going across a period as the effective nuclear charge increases so the bonding electron pair is more strongly attracted by the nucleus.
1Structure - Models of the Particulate of Matter
1.1Introduction to the Particulate Model of Matter
1.2The Nuclear Atom
1.3Electron Configuration
1.4Counting Particles by Mass: The Mole
1.6Elements, Compounds & Mixtures
1.7States of Matter & Changes of State
1.8Reacting Masses &. Volumes
1.9Solutions
2Structure - Models of Bonding & Structure
2.1The Ionic Model
2.2The Covalent Model
2.3Covalent Structures
2.4The Metallic Model
2.5From Models to Materials
2.6Valence Electrons & Ionic Compounds
2.7Molecular Shape
3Structure - Classification of Matter
3.1The Periodic Table: Classification of Elements
3.2Periodic Trends
3.3Group 1 Alkali Metals
3.4Halogens
3.5Noble gases, group 18
3.6Functional Groups: Classification of Organic
3.7Functional Group Chemistry
3.8Alkanes
3.9Alcohols
4Reactivity - What Drives Chemical Reaction?
4.1Endothermic & Exothermic Reactions
4.2Enthalpy of Reaction, Formation, & Hess' Law
5Reactivity - How Much, How Fast & How Far?
5.1Kinetics
5.2Rates of Reaction
5.3Stoichometry
5.4Le Châtelier’s Principle
5.5Introduction to Equilibrium
5.6Equilibrium Constant
5.7Reaction Quotient & Equilibrium Constant
6Reactivity - The Mechanisms of Chemical Change
6.1Proton Transfer Reactions
6.2The pH Scale
6.3Strong & Weak Acids and Bases
6.4Acid Deposition
6.5Types of Organic Reactions
6.6Oxidation & Reduction
6.7Electrochemical Cells
6.9Acid-Base Titrations
6.9.1Titration Calculation Weak Acid & Strong Base
6.9.2Titration Experimental Detail
6.9.3Extended Response - Titration
6.9.4Titration Calculations
6.9.5Titration Curves
6.9.6Titration Calculation Strong Acid & Weak Base
6.9.7IB Multiple Choice - Titrations
6.9.8Polyprotic Acids
6.9.9Titration Calculations Strong Acid & Strong Base
6.9.10Titrations Curves 2
7Measurement, Data Processing & Analysis
7.1Uncertainties & Errors in Measurements & Results
7.2Graphical Techniques
7.3Spectroscopic Identification of Organic Compounds
7.4Infrared Spectroscpy
Jump to other topics
1Structure - Models of the Particulate of Matter
1.1Introduction to the Particulate Model of Matter
1.2The Nuclear Atom
1.3Electron Configuration
1.4Counting Particles by Mass: The Mole
1.6Elements, Compounds & Mixtures
1.7States of Matter & Changes of State
1.8Reacting Masses &. Volumes
1.9Solutions
2Structure - Models of Bonding & Structure
2.1The Ionic Model
2.2The Covalent Model
2.3Covalent Structures
2.4The Metallic Model
2.5From Models to Materials
2.6Valence Electrons & Ionic Compounds
2.7Molecular Shape
3Structure - Classification of Matter
3.1The Periodic Table: Classification of Elements
3.2Periodic Trends
3.3Group 1 Alkali Metals
3.4Halogens
3.5Noble gases, group 18
3.6Functional Groups: Classification of Organic
3.7Functional Group Chemistry
3.8Alkanes
3.9Alcohols
4Reactivity - What Drives Chemical Reaction?
4.1Endothermic & Exothermic Reactions
4.2Enthalpy of Reaction, Formation, & Hess' Law
5Reactivity - How Much, How Fast & How Far?
5.1Kinetics
5.2Rates of Reaction
5.3Stoichometry
5.4Le Châtelier’s Principle
5.5Introduction to Equilibrium
5.6Equilibrium Constant
5.7Reaction Quotient & Equilibrium Constant
6Reactivity - The Mechanisms of Chemical Change
6.1Proton Transfer Reactions
6.2The pH Scale
6.3Strong & Weak Acids and Bases
6.4Acid Deposition
6.5Types of Organic Reactions
6.6Oxidation & Reduction
6.7Electrochemical Cells
6.9Acid-Base Titrations
6.9.1Titration Calculation Weak Acid & Strong Base
6.9.2Titration Experimental Detail
6.9.3Extended Response - Titration
6.9.4Titration Calculations
6.9.5Titration Curves
6.9.6Titration Calculation Strong Acid & Weak Base
6.9.7IB Multiple Choice - Titrations
6.9.8Polyprotic Acids
6.9.9Titration Calculations Strong Acid & Strong Base
6.9.10Titrations Curves 2
7Measurement, Data Processing & Analysis
7.1Uncertainties & Errors in Measurements & Results
7.2Graphical Techniques
7.3Spectroscopic Identification of Organic Compounds
7.4Infrared Spectroscpy
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