ما هي معادلة الأسترة وكيف يتكون الإستر؟ What is the esterification reaction equation?

تفاعل الأسترة هو تفاعل انعكاسي بين حمض كربوكسيلي وكحول في وجود حمض الكبريتيك المركز كعامل حفاز لنزع الماء: R-COOH + R2-OH ⇌ R-COO-R2 + H₂O، ونواتجه إستر وماء. [الصيغة الرمزية الموزونة: R-COOH + R2-OH ⇌ (conc. H₂SO₄) R-COO-R2 + H₂O]

1.4Unit 1. Structure, Bonding and Introduction to Organic Chemistry

Electrons & Ions — IAL Edexcel Linear (Chemistry) Revision Notes & Topical Questions

15 min read 21 parts Lesson 4 of 36 1 quiz

📚 IAL Edexcel Linear (Chemistry) Hub / Topical Revision

Curriculum Equation Key معادلة تفاعل تكوين الإستر (الأسترة) - Esterification Reaction EquationFast Exam Solution

R-COOH + R2-OH ⇌ (conc. H₂SO₄) R-COO-R2 + H₂O

Word Equation & Conditions: حمض كربوكسيلي + كحول (في وجود حمض الكبريتيك المركز كعامل حفاز ونازع للماء) ⇌ إستر + ماء

Location & Specifications: الكيمياء العضوية للثانوية العامة والتحصيلي و Edexcel Chemistry.

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Lesson summary

1.4 Electrons & Ions — IAL EDEXCEL LINEAR (CHEMISTRY) notes: Explain ionisation energy, orbitals and spin, and write electron configurations of atoms and ions.First ionisation energy (IE1) — the energy required to remove one mole of electrons from one mole of gaseous atoms to for…

Second ionisation energy (IE2)
Third ionisation energy (IE3)
Successive ionisation energies
Nuclear charge
Distance from the nucleus
Spin-pair repulsion

Part 1 of 10Ionisation energy definitions

1First ionisation energy (IE1)

First ionisation energy (IE1) the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Measured under standard conditions of 298 K and 101 kPa, in kJ mol^-1.
Ca(g) → Ca⁺(g) + e^-, IE₁ = +590 kJ mol^-1.

2Second ionisation energy (IE2)

Second ionisation energy (IE2) the energy to remove one mole of electrons from one mole of gaseous 1+ ions to form gaseous 2+ ions.
X⁺(g) → X²⁺(g) + e^-, and for calcium IE₂ = 1145 kJ mol^-1.

3Third ionisation energy (IE3)

X²⁺(g) → X³⁺(g) + e^-, and for calcium IE₃ = 4912 kJ mol^-1.

4Successive ionisation energies

Successive ionisation energies each removal needs more energy because electrons leave an increasingly positive ion.

Part 2 of 10Factors affecting ionisation energy

Factors affecting ionisation energy

Nuclear charge

Nuclear charge more protons attract the electrons more strongly.

Distance from the nucleus

Distance from the nucleus attraction falls off rapidly with distance, so outer electrons far from the nucleus are removed more easily.

Shielding

Shielding more inner shells (and subshells) shield the outer electron, making it easier to remove.

Spin-pair repulsion

Spin-pair repulsion two electrons in the same orbital repel, offsetting the nuclear attraction, so paired electrons are removed more easily.

Periodicity

Group 1 metals have low ionisation energies and noble gases have very high ones.

Part 3 of 10Successive ionisation energies and electron configuration

Successive ionisation energies and electron configuration

Uses

Uses predict or confirm configurations, find the number of outer electrons and deduce the group.

Big jumps

Big jumps show a change of shell; smaller jumps show a change of subshell.

Calcium

IE₁ is low because of spin-pair repulsion in the 4s orbital.
IE₂ is higher because there is no spin-pair repulsion.
IE₃ is much higher because the third electron comes from the closer 3p shell.
IE₄ is higher still because the orbital is no longer full and spin-pair repulsion is lower.

Sodium

Sodium a huge jump from IE₁ to IE₂ shows one outer electron, so it is in Group 1 (1s² 2s² 2p⁶ 3s¹).

Magnesium

Magnesium a huge jump from IE₂ to IE₃ shows two outer electrons, so it is in Group 2.

Aluminium

Aluminium a huge jump from IE₃ to IE₄ shows three outer electrons because the fourth electron removed comes from the 2p shell nearer the nucleus.

Deducing the group

Deducing the group the number of electrons removed before the largest jump equals the group number.

Part 4 of 10Shells, subshells and orbitals

Orbital

Orbital a region of space that holds a maximum of two electrons with opposite spins.

Orbitals per subshell

Orbitals per subshell s has 1 (2 electrons), p has 3 (6), d has 5 (10) and f has 7 (14).

Energy

Energy shells rise in energy with principal quantum number n, so n = 4 is above n = 2.
Within a shell the subshells rise s < p < d < f.
A 2p electron has an energy corresponding to n = 2.

Degenerate orbitals

Degenerate orbitals orbitals in the same subshell have equal energy, such as p_x, p_y and p_z.

Exception to the order

3d is slightly higher in energy than 4s, so 4s fills first.

Part 5 of 10Shapes of s and p orbitals

s orbital

s orbital spherical, and it gets larger as the shell number increases.

p orbital

p orbital dumbbell-shaped with two lobes along an axis.
Three p orbitals lie along the x, y and z axes at right angles to each other.
The lobes become larger and longer with increasing shell number.
There is no p subshell in the first shell.

d orbitals

d orbitals their shapes are not required.

Part 6 of 10Electron spin, Pauli principle and Hund's rule

1Electron spin

Electron spin an electron behaves as a tiny spinning charge, clockwise or anticlockwise, creating a tiny magnetic field shown as an arrow up or down.

2Pauli exclusion principle

Pauli exclusion principle an orbital holds a maximum of two electrons and they must have opposite spins.
The energy needed to move to a higher empty orbital is greater than the repulsion between the two electrons, so they pair in the lowest orbital.

3Spin-pair repulsion

Spin-pair repulsion electrons in the same orbital repel each other.

4Hund's rule

Hund's rule electrons occupy separate orbitals of a subshell singly, with parallel spins, before pairing.
Three p electrons occupy p_x, p_y and p_z singly.
A fourth electron pairs up in p_x, leaving two unpaired electrons.

Part 7 of 10Box notation

Boxes

Boxes each box is one atomic orbital, drawn in order of increasing energy from bottom to top.

Arrows

Arrows show the electrons and their spins, with opposite arrows for a pair.

Filling

Filling lowest energy levels fill first, and electrons pair only when no empty orbital remains in the subshell.

Titanium

Titanium the second 4s electron is not promoted to 3d because 3d cannot be half full or full, so it stays in 4s.

Part 8 of 10Writing electron configurations of atoms

Method

Method count the electrons and fill subshells in order of increasing energy: 1s 2s 2p 3s 3p 4s 3d 4p 5s.

Full configuration

Full configuration lists every subshell from 1s upwards.
Potassium (19 electrons) is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹.
Hydrogen is 1s¹.

Shorthand configuration

Shorthand configuration uses the nearest preceding noble gas in square brackets.
Potassium is [Ar] 4s¹, since argon accounts for 18 electrons.
Calcium is [Ar] 4s².

Numerical order

Numerical order write configurations in order of shell, so 3d is written before 4s.
Gallium (31 electrons) is [Ar] 3d¹⁰ 4s² 4p¹.

Part 9 of 10Ions, exceptions and blocks

Positive ions

Positive ions formed by removing electrons from the outer subshell.

Negative ions

Negative ions formed by adding electrons to the outer subshell.

Transition metals

Transition metals fill 4s before 3d but lose 4s electrons first.

Ca²⁺

Ca²⁺ is 1s² 2s² 2p⁶ 3s² 3p⁶, identical to argon.

Exceptions

Exceptions a half-filled or full 3d subshell with one 4s electron is more stable.
Cr is [Ar] 3d⁵ 4s¹, not [Ar] 3d⁴ 4s².
Cu is [Ar] 3d¹⁰ 4s¹, not [Ar] 3d⁹ 4s².

Blocks

Blocks elements are s-, p-, d- or f-block according to the subshell holding the outermost (valence) electron.

Part 10 of 10Electron configuration and chemical properties

Electron configuration and chemical properties

Periodic Table structure

Periodic Table structure periods have the same number of shells and groups have the same number of outer electrons.

Reactivity

Reactivity an element's reactions are determined mainly by its electron configuration.
Elements in one group have similar chemistry.
Across a period the number of outer electrons rises, so chemistry changes from metals to non-metals.

Bonding

Bonding reactions involve losing, gaining or sharing outer electrons to reach a noble-gas arrangement.
Group 1 atoms lose one electron and Group 7 atoms gain one.

Explaining trends

Explaining trends link the electron arrangement and nuclear attraction to the energy needed to remove an electron or to the ease of forming a bond.

Noble gases

He, Ne and Ar have full outer subshells and high ionisation energies.
Their arrangement is very stable, so they rarely gain, lose or share electrons and rarely react.

Key terms

Term	Meaning
First ionisation energy	The energy to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Orbital	A region of space that can hold a maximum of two electrons.
Degenerate	Having the same energy.
Pauli exclusion principle	An orbital holds a maximum of two electrons with opposite spins.
Hund's rule	Electrons occupy orbitals of equal energy singly before pairing.
Spin-pair repulsion	Repulsion between two electrons in the same orbital.
Shorthand configuration	A configuration written with the preceding noble gas in square brackets.
Block	A region of the Periodic Table named after the subshell holding the outermost electron.

Avoid these mistakes

Writing the 4s electrons first in the numerical-order configuration — write 3d before 4s.

Removing 3d electrons before 4s when forming transition-metal ions — 4s electrons are lost first.

Pairing electrons in a p subshell before each orbital has one electron.

Writing Cr as [Ar] 3d⁴ 4s² or Cu as [Ar] 3d⁹ 4s² — both are exceptions.

Explaining a trend by only naming the group — link electron arrangement to ionisation energy or ease of bonding.

Last updated: 18 October 2026