كيف يحسب الرقم الهيدروجيني pH للمحاليل؟ How to calculate pH and pOH?

يحسب الرقم الهيدروجيني بالعلاقة: pH = -log[H⁺]، وللمحاليل المائية عند 25°C: pH + pOH = 14. إذا كان pH 7 فالمحلول قاعدي. [الصيغة الرمزية الموزونة: pH = -log[H⁺] | pOH = -log[OH⁻] | pH + pOH = 14 | [H⁺][OH⁻] = 10⁻¹⁴]

2.1Compound Structure and Properties

Chemical Bonding — AP CHEMISTRY Revision Notes & Topical Questions

5 min read 10 parts Lesson 6 of 13 1 quiz

📚 AP CHEMISTRY Hub / Topical Revision

Curriculum Equation Key قوانين الرقم الهيدروجيني والاتزان الأيوني - pH & Acid-Base Equilibrium FormulasFast Exam Solution

pH = -log[H⁺]  |  pOH = -log[OH⁻]  |  pH + pOH = 14  |  [H⁺][OH⁻] = 10⁻¹⁴

Word Equation & Conditions: الرقم الهيدروجيني = -لوغاريتم تركيز أيون الهيدروجين | pH + pOH = 14

Location & Specifications: الاتزان الأيوني، كيمياء 3 ثانوي، التحصيلي، و AP Chemistry.

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Lesson summary

2.1 Chemical Bonding — AP Notes — AP CHEMISTRY notes: Explain how bond type depends on electronegativity difference and how potential energy curves show bond length and energy.All chemical bonds arise from Coulombic attraction between positive nuclei and negative electrons.A bond…

Dipole moment
Seven elements are diatomic
Bond energy
Bond order
A bond forms because the bonded arrangement has lower potential energy than the separated atoms
Forming a bond releases energy (exothermic); breaking a bond absorbs energy (endothermic)

Part 1 of 8Origin of bonds

Origin of bonds

All chemical bonds arise from Coulombic attraction between positive nuclei and negative electrons.

A bond forms because the bonded arrangement has lower potential energy than the separated atoms.

Forming a bond releases energy (exothermic); breaking a bond absorbs energy (endothermic).

Part 2 of 8Electronegativity and bond polarity

Electronegativity increases across a period (more protons, similar shielding) and decreases down a group (larger radius, more shielding).

Identical atoms, as in H₂ and Cl₂, share electrons equally, giving a nonpolar bond.

Different electronegativities give a polar bond with δ+ and δ− and a bond dipole pointing from δ+ to δ−.

A larger electronegativity difference gives a more polar bond with more ionic character.

Dipole moment

Dipole moment charge × distance (μ = qr), measured in debye (D).

Part 3 of 8Bonding continuum

Large differences, typically metal-nonmetal, give ionic bonding; intermediate differences give polar covalent; very small differences give nonpolar covalent.

AP usually classifies by metal or nonmetal, but ionic compounds have some covalent character and vice versa.

Metals have low ionisation energy and electronegativity, so they lose valence electrons; nonmetals gain or share them.

4Seven elements are diatomic

H₂, N₂, O₂, F₂, Cl₂, Br₂ and I₂ — with nonpolar bonds.

Feature	Ionic	Covalent	Metallic
Particles	Cations and anions	Nonmetal atoms	Metal cations and delocalised electrons
Electrons	Transferred	Shared in localised pairs	Delocalised sea
Typical elements	Metal and nonmetal	Nonmetal and nonmetal	Metals only
Directional	No	Yes	No

Part 4 of 8Properties of ionic and covalent compounds

Properties of ionic and covalent compounds

Ionic compounds are hard, brittle solids with high melting points that conduct only when molten or dissolved, and are often water soluble.

Molecular covalent substances have low melting and boiling points, may be solids, liquids or gases, do not conduct,

and dissolve in water only if polar.

Part 5 of 8Potential energy curves

Far apart, atoms do not interact and the potential energy is zero.

Approaching, nucleus-electron attraction lowers the energy below zero.

At the equilibrium distance the energy is a minimum, and this distance is the bond length.

Closer still, nucleus-nucleus and electron-electron repulsion makes the energy rise steeply.

Bond energy

Bond energy the depth of the well, the energy needed to return to zero potential energy.

A deeper, narrower well represents a stronger, shorter bond.

Part 6 of 8Bond length, order and energy

Bond length, order and energy

Bond length increases with atomic radius, so H–F is stronger than H–Cl, H–Br and H–I.

Higher bond order gives shorter, stronger bonds.

Bond	Order	Length (pm)	Energy (kJ/mol)
C–C	1	154	348
C=C	2	134	614
C≡C	3	120	839

Comparing curves, H₂ has a shorter bond than Cl₂ because H is smaller, and N₂ has a deeper well than F₂ because it has a triple bond.

Part 7 of 8Coulomb's law for ion pairs

Coulomb's law for ion pairs

Potential energy is proportional to q₁q₂/r, so larger charges and smaller separations give stronger attraction.

For ions of the same charge, a larger ion gives a longer equilibrium distance and a shallower well.

So KCl has the shortest distance and deepest well of KCl, KBr and KI.

Part 8 of 8More from the specification

1Electronegativity and nuclear attractive forces

Electronegativity and nuclear attractive forces higher effective nuclear charge exerts stronger attractive forces on bonding electron pairs within covalent bonds.

2Pauling scale

The scale of electronegativity values, with fluorine highest at 4.0.

3Outer shell

Highest-energy occupied shell.

4Attractive force

The pull between particles that holds them together.

Key terms

Term	Meaning
Chemical bond	Electrostatic attraction holding atoms or ions together.
Electronegativity	Ability of an atom in a bond to attract the shared electron pair.
Dipole moment	Charge multiplied by separation, measuring bond or molecule polarity.
Bond length	Internuclear distance at minimum potential energy.
Bond energy	Energy required to break one mole of a bond in the gas phase.
Bond order	Number of shared electron pairs between two atoms.

Avoid these mistakes

Saying electronegativity increases down a group — it decreases

Treating bond energy as negative — it is positive

Reading bond length from the depth of the well — it is the distance at the minimum

Treating the electronegativity cut-offs as sharp boundaries — they are guidelines

Last updated: 18 October 2026