Chemistry Sprint — Expanded Bilingual Study Guide¶

Purpose: A detailed, exam-focused content source for the Chemistry Sprint website. The original study-sheet concepts are preserved, while short explanations, comparisons, worked methods, memory cues, and likely exam prompts have been added for easier revision.

ব্যবহার: এটি Chemistry Sprint ওয়েবসাইটের বিস্তারিত content source। মূল study sheet-এর ধারণা অক্ষুণ্ণ রেখে সহজ ব্যাখ্যা, তুলনা, ধাপে ধাপে হিসাব, মনে রাখার কৌশল ও পরীক্ষাভিত্তিক প্রশ্ন যোগ করা হয়েছে।

How to use this guide / কীভাবে পড়বে¶

Read the Exam-eve gist first. / প্রথমে পরীক্ষার আগের সারাংশ পড়ো।
Memorize the boxed formulas and comparison tables. / সূত্র ও তুলনামূলক টেবিল মনে রাখো।
Practise the worked methods without seeing the answer. / সমাধান না দেখে হিসাবের ধাপ অনুশীলন করো।
Finish with the self-check questions. / শেষে self-check প্রশ্নগুলোর উত্তর দাও।

Chapter 01 — Atomic Structure & Models¶

অধ্যায় ০১ — পরমাণুর গঠন ও মডেল¶

Exam-eve gist / শেষ মুহূর্তের সারাংশ¶

Matter has mass and occupies space. An element cannot be chemically simplified, while a compound contains two or more chemically bonded elements.
An atom retains the identity of an element; a molecule is the smallest independently existing particle of an element or compound.
Electron is negative, proton is positive, and neutron is neutral. Almost all atomic mass is concentrated in the nucleus.
Atomic number, Z = number of protons; mass number, A = protons + neutrons; therefore N = A − Z.
Isotopes have the same Z, isobars have the same A, and isotones have the same neutron number.
Rutherford proved that the atom contains a tiny positive nucleus, but his model could not explain atomic stability or line spectra.
Bohr proposed stationary energy levels. Radiation is emitted or absorbed only when an electron jumps between levels: ΔE = hν.
Four quantum numbers specify an electron. Electron filling follows Aufbau, Pauli and Hund.

1. Matter, element, compound, atom and molecule¶

Term	English explanation	বাংলা ব্যাখ্যা	Example
Matter	Anything that has mass and occupies space	যার ভর আছে ও স্থান দখল করে	Air, water, iron
Element	Pure substance that cannot be chemically broken into simpler substances	রাসায়নিকভাবে সরল পদার্থে ভাঙা যায় না	H, O, Fe
Compound	Two or more elements chemically bonded in a fixed ratio	দুই বা ততোধিক মৌল নির্দিষ্ট অনুপাতে যুক্ত	H₂O, CO₂
Atom	Smallest particle retaining an element's identity	মৌলের ধর্ম বহনকারী ক্ষুদ্রতম কণা	H atom, Mg atom
Molecule	Smallest independent particle of an element or compound	মৌল বা যৌগের ক্ষুদ্রতম স্বাধীন কণা	O₂, H₂O

2. Fundamental particles¶

Particle	Symbol	Charge	Approximate mass	Position
Electron	e⁻	−1	`9.11 × 10⁻²⁸ g`	Outside nucleus
Proton	p⁺	+1	`1.673 × 10⁻²⁴ g ≈ 1 amu`	Nucleus
Neutron	n⁰	0	`1.675 × 10⁻²⁴ g ≈ 1 amu`	Nucleus

The nucleus occupies a tiny fraction of atomic volume but contains nearly all its mass. A neutral atom has equal numbers of protons and electrons. When it loses electrons it becomes a cation; when it gains electrons it becomes an anion.

নিউক্লিয়াস আয়তনে খুব ছোট হলেও পরমাণুর প্রায় সম্পূর্ণ ভর বহন করে। নিরপেক্ষ পরমাণুতে proton ও electron সমান। Electron হারালে cation এবং গ্রহণ করলে anion হয়।

3. Atomic number, mass number and relative atomic mass¶

A = Z + N

Z identifies the element and equals the number of protons.
A is a whole number for one isotope.
Relative atomic mass may be fractional because it is the abundance-weighted average of naturally occurring isotopes.
1 amu = 1/12 of the mass of one carbon-12 atom ≈ 1.6605 × 10⁻²⁴ g.

Weighted-average method

Average atomic mass = Σ(isotopic mass × percentage abundance) / 100

Example: if X has 75% X-35 and 25% X-37,

Average mass = (35 × 75 + 37 × 25)/100 = 35.5 amu

4. Isotope, isobar and isotone¶

Relation	Same	Different	Example idea
Isotope	Atomic number `Z`	Mass number `A`	¹H, ²H, ³H
Isobar	Mass number `A`	Atomic number `Z`	⁴⁰Ar, ⁴⁰Ca
Isotone	Neutron number `N`	`Z` and `A`	¹⁴C and ¹⁵N

Memory cue: same Z → isotope; same A → isobar; same N → isotone.

5. Radioactivity and radiation¶

Radiation	Nature	Charge	Penetration	Ionisation
α	Helium nucleus, ⁴₂He²⁺	+2	Low	Very high
β	Fast electron stream	−1	Medium	Medium
γ	Electromagnetic radiation	0	Very high	Low

Important uses include Co-60 in cancer treatment, I-131 for thyroid diagnosis/treatment, C-14 for age determination and P-32 as a plant tracer. Radiation must be handled with shielding, distance and limited exposure time.

6. Rutherford model¶

From α-particle scattering Rutherford concluded:

Most α particles passed straight through → most of the atom is empty space.
A few were deflected → positive charge is concentrated.
Very few rebounded → nucleus is extremely small, dense and massive.

Limitations: An orbiting charged electron should continuously radiate energy and fall into the nucleus. The model also could not explain discrete line spectra or electron arrangement.

7. Bohr model¶

Electrons move only in permitted stationary orbits without losing energy.
Angular momentum is quantized: mvr = nh/2π.
Radiation appears only during a transition: ΔE = E₂ − E₁ = hν.
A downward jump emits energy; an upward jump absorbs energy.

Bohr explains hydrogen and hydrogen-like ions well, but not fine spectra, Zeeman/Stark effects or complex multi-electron atoms.

8. Quantum numbers¶

Quantum number	Symbol	Meaning	Allowed values
Principal	`n`	Shell, energy and size	1, 2, 3…
Azimuthal	`l`	Subshell/shape	`0` to `n−1`; s=0, p=1, d=2, f=3
Magnetic	`mₗ`	Orbital orientation	`−l` to `+l`
Spin	`mₛ`	Electron spin	`+½` or `−½`

Subshell capacities: s², p⁶, d¹⁰, f¹⁴; maximum shell capacity: 2n².

9. Electron-filling rules¶

Aufbau: lower-energy orbitals fill first. Lower n+l fills first; if equal, lower n fills first.
Pauli exclusion: one orbital holds at most two electrons with opposite spin; no two electrons in an atom have identical four quantum numbers.
Hund: degenerate orbitals fill singly with parallel spin before pairing.
Common order: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p…
Chromium and copper show stable half/full-filled exceptions: Cr [Ar]3d⁵4s¹; Cu [Ar]3d¹⁰4s¹.

Common mistakes / সাধারণ ভুল¶

Do not confuse mass number with average atomic mass.
Proton charge is positive, even if a source typo shows a negative sign.
Isotopes have nearly identical chemical behaviour but different masses and some physical properties.
Bohr orbit and quantum-mechanical orbital are not the same concept.

Self-check / নিজেকে যাচাই¶

Why is relative atomic mass often fractional?
Differentiate isotope, isobar and isotone with one example each.
Explain Rutherford's observations and conclusions.
State Bohr's postulates and two limitations.
Write the four quantum numbers of the last electron of a selected atom.

Chapter 02 — Periodic Table & Periodic Properties¶

অধ্যায় ০২ — পর্যায় সারণি ও পর্যায়বৃত্ত ধর্ম¶

Exam-eve gist¶

Newlands used octaves, Mendeleev arranged elements mainly by atomic mass, and Moseley established atomic number as the modern basis.
The long periodic table contains 7 periods and 18 groups.
The last-filled subshell identifies the s, p, d or f block.
Across a period: effective nuclear charge generally rises, radius falls, while ionization energy and electronegativity rise.
Down a group: new shells increase size; ionization energy and electronegativity generally fall.
Stable configurations cause important exceptions such as IE(Be) > IE(B) and IE(N) > IE(O).

1. Development of the table¶

Scientist	Basis/contribution	Main limitation or importance
Newlands	Every eighth element showed similarity when ordered by mass	Worked mainly for lighter elements
Mendeleev	Properties periodic with atomic mass	Left gaps and predicted undiscovered elements, but mass order had anomalies
Moseley	Properties periodic with atomic number	Basis of the modern periodic law

Modern periodic law: Physical and chemical properties of elements are periodic functions of their atomic numbers.

Ionization Energy: Extracting/pulling an electron away from a neutral atom to form a cation (positive ion).
Electron Affinity: Adding/pushing an electron into a neutral atom to form an anion (negative ion).
Electronegativity: The tug-of-war where an atom pulls shared electrons toward itself inside a covalent bond.

That is the exact mental framework you need for your exams.

2. Major families¶

Group 1: alkali metals, ns¹, highly reactive, usually +1.
Group 2: alkaline-earth metals, ns², usually +2.
Group 17: halogens, ns²np⁵, reactive nonmetals, usually −1.
Group 18: noble gases, filled shell ns²np⁶ except He=1s²; monatomic and generally inert.
The staircase region contains metalloids such as B, Si, Ge, As, Sb and Te.

Noble gases have low melting/boiling points and are colourless, odourless and nonflammable. Uses include He in balloons/cryogenics, Ne in signs, Ar in bulbs and welding, and Xe in specialised lamps.

3. Blocks and electronic configuration¶

Block	General ending	Location
s	`ns¹–²`	Groups 1–2, plus He by configuration
p	`ns²np¹–⁶`	Groups 13–18
d	`(n−1)d¹–¹⁰ns⁰–²`	Groups 3–12
f	`(n−2)f¹–¹⁴`	Lanthanides and actinides

Period number is generally the highest principal quantum number. For p-block elements, group = 10 + valence electrons; for many d-block elements, group = (n−1)d electrons + ns electrons.

4. Radius and effective nuclear charge¶

Zeff = Z − S

Across a period, shielding changes only modestly while nuclear charge rises, so electrons are pulled inward. Down a group, added shells and shielding dominate, so radius increases.

cation < parent atom < anion

In an isoelectronic series, the species with more protons is smaller. Example: O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺.

Types of radius:

Covalent radius = half the distance between identical covalently bonded nuclei.
Metallic radius = half the distance between neighbouring metal nuclei.
van der Waals radius describes closest nonbonded approach and is usually larger.

5. Ionization energy (IE)¶

Ionization energy is the minimum energy needed to remove the most loosely held electron from an isolated gaseous atom.

Factors: atomic size, nuclear charge, shielding, penetration, and stable half/full-filled subshells. Successive ionization energies always rise; a large jump indicates removal from a stable inner shell and helps reveal valency.

Important exceptions:

IE(Be) > IE(B) because Be has stable 2s², while B loses a higher-energy 2p¹ electron.
IE(N) > IE(O) because N has stable half-filled 2p³; O contains one paired p electron with extra repulsion.

6. Electronegativity and electron affinity¶

Electronegativity is an atom's attraction for the shared electron pair in a bond. Fluorine is highest. A larger electronegativity difference generally increases ionic character.

Electron affinity is the energy change when an isolated gaseous atom accepts an electron. It generally becomes more favourable across a period, but stable configurations and very compact orbitals cause exceptions. EA(Cl) is more favourable than EA(F) because compact fluorine has stronger electron-electron repulsion.

7. Trend map¶

Property	Across a period →	Down a group ↓
Atomic radius	Decreases	Increases
Metallic character	Decreases	Increases
Ionization energy	Generally increases	Generally decreases
Electronegativity	Generally increases	Generally decreases
Effective nuclear charge	Increases	Changes less regularly

Common mistakes¶

Do not claim every d-block element is a transition element; Zn, Cd and Hg have filled d¹⁰ ions/atoms in common states.
Electron affinity and electronegativity are related but not identical.
Trends are general patterns, not exception-free laws.

Self-check¶

Why did atomic number solve Mendeleev's Ar/K anomaly?
Find the period, group and block from a given electronic configuration.
Arrange an isoelectronic series by size and justify it.
Explain the Be/B, N/O and Cl/F exceptions.

Chapter 03 — Chemical Bonds¶

অধ্যায় ০৩ — রাসায়নিক বন্ধন¶

Exam-eve gist¶

Atoms bond to achieve lower energy and often a stable noble-gas-like valence shell.
Complete electron transfer forms ionic bonds; sharing forms covalent bonds; donation of both shared electrons by one atom forms a coordinate bond.
Delocalized electrons explain metallic conductivity, lustre, malleability and ductility.
Hydrogen bonding strongly affects boiling point, solubility and the open structure of ice.
Lewis structures count valence electrons; resonance means the real molecule is a delocalized hybrid.

1. Octet, duplet and valency¶

The octet rule says many main-group atoms combine so that the valence shell reaches ns²np⁶. Hydrogen and lithium are commonly described through a two-electron duplet. Valency is the number of electrons lost, gained or shared to achieve a stable arrangement.

The octet rule is a useful model, not a universal law. BF₃ is electron-deficient, PCl₅ has an expanded valence shell, and odd-electron species also exist.

2. Types of bond¶

Bond	Electron behaviour	Typical participants	Example
Ionic	Complete transfer and electrostatic attraction	Metal + nonmetal	NaCl, MgO
Covalent	Pair(s) shared	Nonmetal + nonmetal	H₂, O₂, CH₄
Coordinate	Both bonding electrons supplied by one donor	Lone-pair donor + acceptor	NH₄⁺, H₃O⁺, BF₃←NH₃
Metallic	Positive cores in delocalized electron sea	Metal atoms	Cu, Fe, Al

3. Formation and properties of ionic compounds¶

For NaCl, Na loses one electron to become Na⁺ and Cl gains it to become Cl⁻. The lattice is held by strong nondirectional electrostatic forces.

High melting/boiling points due to strong lattice attraction.
Brittle: shifting layers can bring like charges together and split the crystal.
Conduct electricity when molten or dissolved because ions can move, but not in the solid state.
Commonly soluble in polar solvents such as water.

4. Covalent bonds¶

One, two or three shared pairs form single, double or triple bonds. Increasing bond order usually shortens and strengthens a bond.

Molecular covalent substances often have low melting points and poor conductivity.
Giant covalent networks such as diamond have very high melting points.
Graphite conducts along layers because it has delocalized electrons.
Bond polarity arises when atoms have unequal electronegativity.

5. Coordinate bonding¶

A Lewis base donates a lone pair and a Lewis acid accepts it. After formation, a coordinate covalent bond behaves like an ordinary covalent bond.

Examples:

NH₃ + H⁺ → NH₄⁺
H₂O + H⁺ → H₃O⁺
BF₃ + :NH₃ → F₃B←NH₃

6. Metallic and intermolecular forces¶

Metallic bonding is nondirectional. Mobile delocalized electrons carry charge and heat, while metal-ion layers can slide without destroying the attraction.

Intermolecular forces are weaker than primary bonds:

London/van der Waals forces occur in all particles and grow with polarizability.
Dipole–dipole attraction occurs between polar molecules.
Hydrogen bonding occurs when H bonded to F, O or N is attracted to a lone pair on F/O/N.

Water forms an open tetrahedral H-bonded arrangement in ice. The open spaces lower the density, which is why ice floats.

7. Lewis structures and resonance¶

Method:

Count total valence electrons; adjust for charge.
Select a central atom, usually the least electronegative except H.
Connect atoms with single bonds.
Complete terminal octets, then place remaining electrons on the centre.
If the centre lacks an octet, form multiple bonds where reasonable.
Check formal charges and choose structures with sensible charge distribution.

Resonance contributors are alternative drawings, not molecules switching back and forth. The actual structure is a lower-energy resonance hybrid; benzene's equal C–C bonds are a classic example.

Common mistakes¶

Ionic compounds contain formula units, not discrete NaCl molecules in the crystal.
Coordinate bond origin differs, but the completed bond is not permanently a different physical bond type.
Hydrogen bond is much weaker than O–H or N–H covalent bonds.

Self-check¶

Explain NaCl formation with electron configuration.
Compare ionic and covalent compounds in a table.
Why does ice float?
Draw Lewis structures for CO₂, NH₄⁺ and BF₃ and discuss octet behaviour.

Chapter 04 — Defects in Solids & Semiconductors¶

অধ্যায় ০৪ — কঠিনের ত্রুটি ও অর্ধপরিবাহী¶

Exam-eve gist¶

Real crystals contain irregularities called defects; point defects are localized, while line defects extend along a row.
Stoichiometric defects preserve the compound ratio; non-stoichiometric defects change it; impurity defects arise from foreign atoms/ions.
Schottky defect creates paired vacancies and lowers density. Frenkel defect moves an ion to an interstitial site and usually leaves density unchanged.
Conductors have overlapping/partially filled bands, insulators have a large band gap, and semiconductors have a small band gap.
Group 15 doping of Si/Ge produces n-type; group 13 doping produces p-type.

1. Crystal classification¶

Crystal	Main force	Typical property	Example
Covalent/atomic	Strong covalent network	Hard, high melting	Diamond, quartz
Ionic	Electrostatic attraction	Brittle, high melting	NaCl
Metallic	Metallic bonding	Conductive, malleable	Cu, Fe
Molecular	Intermolecular force	Soft, lower melting	Ice, dry ice

2. Defect hierarchy¶

Point defects: vacancy, interstitial, substitutional impurity, Schottky and Frenkel.
Line defects: edge or screw dislocation along a line of lattice sites.
Stoichiometric/intrinsic defects: chemical formula ratio remains unchanged.
Non-stoichiometric defects: metal excess or metal deficiency changes ideal ratio.
Impurity defects: deliberate or accidental foreign species occupy sites.

3. Schottky vs Frenkel¶

Feature	Schottky	Frenkel
What happens	Equal numbers of cations and anions leave lattice sites	Usually a smaller ion leaves its site and occupies an interstitial
Electrical neutrality	Preserved	Preserved
Stoichiometry	Preserved	Preserved
Density	Decreases	Nearly unchanged
Favoured by	Similar ion sizes, high coordination	Large size difference, low coordination
Examples	NaCl, KCl, CsCl	AgCl, AgBr, ZnS

4. Non-stoichiometric defects¶

Metal excess:

Anion vacancies may trap electrons; these are F-centres and can give colour.
Extra cations may occupy interstitial sites with electrons maintaining neutrality.

Metal deficiency: Some cations are missing and nearby metal ions adopt a higher oxidation state to preserve charge balance. This is common in compounds of variable-valency metals.

5. Energy-band theory¶

When many atoms form a solid, closely spaced atomic orbitals create bands.

Valence band: highest band normally occupied by electrons.
Conduction band: band in which electrons can move through the solid.
Band gap (Eg): forbidden energy gap between them.

Material	Band picture	Conductivity
Conductor	Bands overlap or a band is partially filled	High
Semiconductor	Small `Eg`; heat/light can promote electrons	Moderate, rises with temperature
Insulator	Large `Eg`	Very low

6. Intrinsic and extrinsic semiconductor¶

Pure Si or Ge is an intrinsic semiconductor. Thermal energy creates an electron in the conduction band and a corresponding hole in the valence band; electrons and holes are equal in number.

Doping produces extrinsic semiconductors:

Type	Dopant	Majority carrier	Minority carrier
n-type	Pentavalent Group 15, e.g. P/As/Sb	Electron	Hole
p-type	Trivalent Group 13, e.g. B/Al/Ga	Hole	Electron

The crystal remains electrically neutral overall. “p-type” does not mean the material has a net positive charge, and “n-type” does not mean a net negative charge.

7. Why defects matter¶

Defects can change density, colour, diffusion, mechanical strength, chemical reactivity and electrical conductivity. Controlled impurity is the foundation of semiconductor devices, while unwanted defects can weaken a material.

Common mistakes¶

Vacancy/interstitial are geometrical descriptions; Schottky/Frenkel specify charge-balanced defect patterns in ionic solids.
Do not say Frenkel lowers density significantly; the ion remains inside the crystal.
Holes are effective positive charge carriers, not independent positive particles like protons.

Self-check¶

Draw a hierarchy of solid defects.
Compare Schottky and Frenkel defects.
Explain an F-centre and its connection to colour.
Compare conductor, semiconductor and insulator by band theory.
Explain p-type and n-type doping without implying net crystal charge.

Chapter 05 — Chemical Equilibrium & Thermochemistry¶

অধ্যায় ০৫ — রাসায়নিক সাম্য ও তাপ-রসায়ন¶

Exam-eve gist¶

Equilibrium is dynamic: forward and reverse reactions continue at equal rates while macroscopic concentrations stay constant.
For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d/[A]^a[B]^b.
Pure solids and pure liquids are omitted from equilibrium expressions.
Kp = Kc(RT)^Δn, where Δn = gaseous product moles − gaseous reactant moles.
A catalyst speeds both directions equally and does not change K or the equilibrium composition.
Temperature changes K; concentration and pressure shift position but do not change K at fixed temperature.
Hess's law works because enthalpy is a state function.

1. Reversible, irreversible and dynamic equilibrium¶

An irreversible reaction proceeds mainly in one direction (→). A reversible reaction can proceed in both directions (⇌). In a closed system, a reversible reaction may reach equilibrium where:

forward rate = reverse rate;
reactant and product concentrations are constant, not necessarily equal;
equilibrium can be approached from either side;
the reaction is usually incomplete;
external conditions determine the equilibrium position.

2. Equilibrium constants¶

For aA(g) + bB(g) ⇌ cC(g) + dD(g):

Kc = ([C]^c[D]^d)/([A]^a[B]^b)

Kp = (P_C^c P_D^d)/(P_A^a P_B^b)

The exponents come from the balanced equation. Pure solids and liquids have effectively constant activity and are omitted. A large K means products are favoured at equilibrium; a small K means reactants are favoured. It says nothing directly about reaction speed.

3. Relationship between Kp and Kc¶

Kp = Kc(RT)^Δn

Only gaseous coefficients count in Δn. If Δn = 0, then Kp = Kc.

Method: balance equation → count gaseous product moles → subtract gaseous reactant moles → substitute R and absolute temperature T.

4. Reaction quotient and direction¶

Q has the same form as K but uses the current, not necessarily equilibrium, composition.

Q < K: reaction proceeds forward.
Q > K: reaction proceeds backward.
Q = K: system is at equilibrium.

5. Le Chatelier's principle¶

A system shifts in the direction that reduces an imposed change.

Change	General response
Add reactant	Shift toward products
Remove reactant	Shift toward reactants
Add product	Shift toward reactants
Remove product	Shift toward products
Increase pressure	Shift to fewer gaseous moles
Decrease pressure	Shift to more gaseous moles
Add catalyst	No shift; equilibrium reached faster

Pressure has no effect when gaseous moles are equal on both sides. Inert gas at constant volume does not change reacting-gas partial pressures, so it does not shift equilibrium.

6. Temperature and heat¶

Treat heat as a reactant in an endothermic reaction and as a product in an exothermic reaction.

Increasing temperature favours the endothermic direction.
Decreasing temperature favours the exothermic direction.
Temperature is the factor that changes the numerical value of K.

7. Enthalpy and thermochemical equations¶

Endothermic: absorbs heat, ΔH > 0.
Exothermic: releases heat, ΔH < 0.
Reaction enthalpy: ΔH = H(products) − H(reactants).

A thermochemical equation must show balanced coefficients, physical states, temperature/conditions when relevant, and ΔH for the equation exactly as written.

8. Hess's law¶

The total enthalpy change is independent of path because enthalpy is a state function.

Rules:

Reverse an equation → reverse the sign of ΔH.
Multiply coefficients by n → multiply ΔH by n.
Add equations → add their ΔH values.
Cancel intermediates to obtain the target equation.

Common mistakes¶

Equal rates do not mean equal concentrations.
Catalyst does not increase equilibrium yield.
Do not include solids or pure liquids in Kc.
Use Kelvin, not Celsius, in Kp = Kc(RT)^Δn.
When manipulating a Hess equation, manipulate ΔH identically.

Self-check¶

Why is equilibrium called dynamic?
Write Kc and Kp for a supplied balanced equation.
Predict the effect of pressure, temperature and catalyst on the Haber process.
Solve a Hess-law problem by equation manipulation.

Chapter 06 — Chemical Kinetics¶

অধ্যায় ০৬ — রাসায়নিক গতিবিদ্যা¶

Exam-eve gist¶

Rate is concentration change per unit time: reactant rate is written negative because its concentration falls.
Experimental rate law: rate = k[A]^x[B]^y; overall order is x+y.
Molecularity belongs to an elementary step and is a positive integer; order is experimental and may be zero, fractional or integer.
Zero order: x = kt; first order: k = (2.303/t)log(a/(a−x)).
First-order half-life is t½ = 0.693/k and does not depend on initial concentration.
A reaction becomes pseudo-first-order when one reactant is present in such large excess that its concentration remains effectively constant.

1. Rate of reaction¶

For aA → pP:

Rate = −(1/a)d[A]/dt = (1/p)d[P]/dt

The stoichiometric division makes the calculated rate independent of which species is monitored. Common unit: mol L⁻¹ s⁻¹.

Rate is affected by concentration/pressure, temperature, catalyst, surface area, reactant nature and sometimes light.

2. Rate law and rate constant¶

rate = k[A]^x[B]^y

The exponents are found experimentally and do not automatically equal balanced-equation coefficients, except in a truly elementary step. k depends on temperature and catalyst; its unit depends on reaction order.

For overall order n, unit of k is (concentration)^(1−n) time⁻¹.

3. Molecularity vs order¶

Feature	Molecularity	Order
Meaning	Particles in one elementary step	Sum of experimental rate-law powers
Possible values	Positive whole number, usually 1–3	Zero, integer or fractional
Applies to	Elementary step	Overall reaction/rate law
Determination	Mechanism	Experiment

4. Zero-order reaction¶

rate = k

If initial concentration is a and amount reacted is x after time t:

x = kt, or remaining concentration [A]t = [A]0 − kt.

Plot of [A] vs t is a straight line with slope −k.
Unit of k: concentration/time.
Half-life: t½ = [A]0/(2k); it depends on initial concentration.

5. First-order reaction¶

rate = k[A]

k = (2.303/t)log(a/(a−x)) = (2.303/t)log([A]0/[A]t)

Plot of log[A] vs t is straight with slope −k/2.303.
Unit of k: time⁻¹.
The reaction approaches completion asymptotically.

For half-life, set [A]t = [A]0/2:

t½ = (2.303/k)log2 = 0.693/k

Example: if t½ = 20 min, k = 0.693/20 = 0.03465 min⁻¹.

6. Pseudo-first-order reaction¶

For rate = k[A][B], if B is in huge excess, [B] changes negligibly:

rate = k[B][A] = k′[A]

Thus a higher-molecularity reaction behaves kinetically like first order. Sucrose hydrolysis in excess water is a standard example.

7. Numerical checklist¶

Identify order and correct integrated equation.
Keep time units consistent with the desired k unit.
Distinguish amount reacted x from amount remaining a−x.
Use ordinary log with 2.303, or natural log without it.
Check dimensions and significant figures.

Common mistakes¶

Order cannot usually be read directly from a balanced overall equation.
First-order half-life is concentration-independent; zero-order half-life is not.
Do not insert x where the formula requires a−x.

Self-check¶

Compare order and molecularity.
Derive the first-order half-life equation.
Determine k from concentration-time data.
Explain why sucrose hydrolysis can be pseudo-first-order.

Chapter 07 — pH & Solution Concentration¶

অধ্যায় ০৭ — pH ও দ্রবণের ঘনমাত্রা¶

Exam-eve gist¶

pH = −log[H⁺], pOH = −log[OH⁻].
At 25°C, Kw = [H⁺][OH⁻] = 10⁻¹⁴ and pH + pOH = 14.
Lower pH means higher hydrogen-ion concentration; one pH unit represents a tenfold change.
Molarity uses litres of solution, molality uses kilograms of solvent, and normality uses equivalents per litre of solution.
N = M × n-factor for a specified reaction.
Electrolyte conductivity depends on ion concentration, mobility, temperature, solvent and pH.

1. Meaning and scale of pH¶

At 25°C:

pH < 7: acidic.
pH = 7: neutral pure water.
pH > 7: basic.

The scale is logarithmic. A solution of pH 3 has ten times the [H⁺] of pH 4 and one hundred times that of pH 5. Typical examples from the source include stomach medium near pH 2, arterial blood near pH 7.4, skin commonly mildly acidic, and soils over a broad approximate range.

2. Ionic product of water¶

Water undergoes slight self-ionization:

H₂O ⇌ H⁺ + OH⁻

Because liquid-water concentration is effectively constant:

Kw = [H⁺][OH⁻]

At 25°C, Kw = 1.0 × 10⁻¹⁴. In pure water [H⁺]=[OH⁻], so each is 10⁻⁷ M, giving pH 7.

Taking negative logarithms:

pKw = pH + pOH, and at 25°C, pH + pOH = 14.

3. Fast pH calculation method¶

Write complete dissociation for a strong acid/base.
Determine [H⁺] or [OH⁻], including stoichiometric multiplier.
Apply −log.
If starting from pOH, calculate pH = 14 − pOH at 25°C.

Examples:

0.01 M HCl → [H⁺]=10⁻² M → pH=2.
Using the study-sheet simplification, 0.005 M H₂SO₄ → [H⁺]=0.010 M → pH=2.
If pH=5.50, then pOH=8.50 and [OH⁻]=10⁻⁸·⁵ M.

4. Mole and concentration measures¶

number of moles = mass / molar mass

Measure	Formula	Denominator	Temperature effect
Molarity, M	moles solute / L solution	Volume of total solution	Changes with temperature
Molality, m	moles solute / kg solvent	Mass of solvent only	Essentially independent
Normality, N	equivalents / L solution	Volume of total solution	Changes with temperature

N = M × n-factor

The n-factor depends on the reaction. For acid-base reactions it often reflects replaceable H⁺ or OH⁻; for redox reactions it reflects electrons exchanged. Therefore normality is reaction-specific.

5. Dilution and preparation¶

For dilution where solute amount stays constant:

M₁V₁ = M₂V₂

To prepare a molar solution: calculate required moles → convert to mass → dissolve in some solvent → transfer to volumetric flask → make up exactly to the mark → mix thoroughly.

Never confuse “1 L solvent” with “solution made up to 1 L.”

6. Electrical properties¶

Electrolytes conduct because mobile ions carry charge. Conductivity generally rises with the number and mobility of ions, though dilution effects must distinguish specific conductivity from molar conductivity.

Electrophoresis: charged colloidal particles move toward the oppositely charged electrode.
Electro-osmosis: the dispersion medium moves while charged particles are relatively fixed.
A conductivity meter measures a solution's ability to carry current.

Common mistakes¶

The pH 0–14 range is the common dilute-aqueous scale at ordinary conditions, not an absolute universal limit.
pH=7 is neutral specifically at about 25°C; neutrality always means [H⁺]=[OH⁻].
Molarity uses solution volume; molality uses solvent mass.
Always state the reaction when using normality.

Self-check¶

Derive pH+pOH=14 from Kw.
Calculate pH from a strong-acid concentration and reverse-calculate concentration from pH.
Compare M, m and N with units.
Explain electrophoresis and electro-osmosis.

Chapter 08 — Organic Chemistry Reactions¶

অধ্যায় ০৮ — জৈব রসায়নের বিক্রিয়া¶

Exam-eve gist¶

Grignard reagent RMgX is prepared from an alkyl/aryl halide and Mg in dry ether; it is destroyed by water and forms new C–C bonds.
Methanal + Grignard gives a primary alcohol; another aldehyde gives secondary; ketone gives tertiary after acidic work-up.
Aldehydes normally give Tollens silver mirror; aliphatic aldehydes give Fehling brick-red Cu₂O, while ordinary ketones do not.
Primary amide gives a one-carbon-shorter primary amine by Hofmann degradation.
Carbonyl with α-H → aldol; aldehyde without α-H → Cannizzaro; methyl carbonyl → haloform.
Benzene commonly undergoes electrophilic substitution. Ordinary HX addition to an unsymmetrical alkene follows Markovnikov orientation.

1. Grignard reagent¶

Preparation:

R–X + Mg —dry ether→ R–MgX

The C–Mg bond is strongly polarized, so carbon behaves nucleophilically. Moisture or any acidic hydrogen destroys the reagent; all apparatus and ether must therefore be dry.

Main applications after hydrolysis/work-up:

Reactant with RMgX	Product
Water/proton source	Hydrocarbon, RH
CO₂	Carboxylic acid, RCOOH
Methanal	Primary alcohol
Other aldehyde	Secondary alcohol
Ketone	Tertiary alcohol

Grignard chemistry is valuable because it creates a new carbon-carbon bond and lengthens the carbon skeleton.

2. Functional groups¶

Class	General group	Example
Alcohol	`R–OH`	CH₃CH₂OH
Aldehyde	`R–CHO`	CH₃CHO
Ketone	`R–CO–R′`	CH₃COCH₃
Amine	`R–NH₂`, `R₂NH`, `R₃N`	CH₃NH₂
Amide	`R–CONH₂` etc.	CH₃CONH₂

Alcohols are primary, secondary or tertiary according to how many carbon groups attach to the carbon bearing OH. Aldehydes have at least one H on the carbonyl carbon; ketones have two carbon groups.

3. Aldehyde vs ketone tests¶

Test	Aldehyde result	Ordinary ketone result
Tollens reagent, `[Ag(NH₃)₂]⁺`	Silver mirror/Ag deposit	No reaction
Fehling solution	Aliphatic aldehyde gives brick-red Cu₂O	No reaction

Aldehydes are readily oxidized to carboxylic acids, while ketones resist these mild oxidants. Clean glassware is important for a proper Tollens mirror, and Tollens reagent should be freshly prepared and safely disposed of.

4. Amines, amides and Hofmann degradation¶

Amines are ammonia derivatives; amides contain a carbonyl directly attached to nitrogen. In Hofmann bromamide degradation:

RCONH₂ + Br₂ + 4KOH → RNH₂ + 2KBr + K₂CO₃ + 2H₂O

The product is a primary amine with one fewer carbon atom than the starting amide.

5. Aldol condensation¶

Aldehydes or ketones containing at least one α-hydrogen form an enolate/enol in dilute base. It attacks another carbonyl molecule to give a β-hydroxy carbonyl compound (aldol), which may dehydrate to an α,β-unsaturated carbonyl compound.

Example: two ethanal molecules form 3-hydroxybutanal, followed by dehydration to crotonaldehyde under suitable conditions.

6. Cannizzaro reaction¶

Aldehydes without α-hydrogen undergo disproportionation in concentrated alkali: one molecule is oxidized to carboxylate while another is reduced to alcohol.

2HCHO + NaOH → HCOONa + CH₃OH

Decision cue: carbonyl has α-H → consider aldol; aldehyde lacks α-H → consider Cannizzaro.

7. Haloform reaction¶

Methyl ketones containing CH₃CO–, and compounds oxidizable to that group, react with halogen in base. With iodine, yellow iodoform precipitate forms:

CH₃CO–R + I₂/OH⁻ → CHI₃↓ + carboxylate

The iodoform test is therefore used for methyl ketones and certain secondary alcohols/ethanol that produce methyl carbonyl compounds on oxidation.

8. Electrophile, nucleophile, carbocation and carbanion¶

Species	Electron behaviour	Typical feature/example
Electrophile	Accepts an electron pair	H⁺, NO₂⁺, carbocation
Nucleophile	Donates an electron pair	OH⁻, CN⁻, NH₃, RMgX
Carbocation	Positively charged carbon	Electron-deficient, usually planar
Carbanion	Negatively charged carbon	Electron-rich, nucleophilic

Carbocation stability commonly follows 3° > 2° > 1° > methyl due to hyperconjugation and inductive donation; resonance-stabilized allylic/benzylic cases need separate comparison. Simple alkyl carbanions often show the reverse trend.

9. Aromatic electrophilic substitution¶

Nitration: concentrated HNO₃/H₂SO₄ generates NO₂⁺. Benzene attacks the electrophile, forms a σ-complex, then loses H⁺ to restore aromaticity. –NO₂ is deactivating and meta-directing.

Friedel–Crafts: an alkyl or acyl electrophile is generated using a Lewis acid such as AlCl₃. Alkylation can rearrange and over-alkylate; acylation generally avoids these problems but requires later reduction if an alkylbenzene is desired.

10. Addition and halogenation¶

For ordinary addition of HX to an unsymmetrical alkene, Markovnikov's rule predicts H attaches to the carbon already bearing more H atoms, while X attaches to the more substituted carbon. The orientation follows formation of the more stable carbocation intermediate. The peroxide effect is a special radical exception mainly associated with HBr.

Alkenes add Br₂/Cl₂ across the π bond and decolourize bromine solution.
Alkanes undergo radical halogenation with light or heat.
Benzene undergoes substitution with halogen in the presence of FeCl₃/FeBr₃ or AlCl₃-type catalyst.

11. Reaction-selection map¶

Clue in question	Likely reaction/test
`R–X + Mg`, dry ether	Grignard preparation
Aldehyde identification	Tollens/Fehling
Primary amide + Br₂/KOH	Hofmann degradation
Carbonyl compound with α-H + dilute base	Aldol
Aldehyde without α-H + concentrated base	Cannizzaro
`CH₃CO–` or suitable precursor + I₂/OH⁻	Iodoform
Benzene + HNO₃/H₂SO₄	Nitration
Alkene + HX	Markovnikov addition unless exception stated

Common mistakes¶

Grignard reactions require dry conditions; water is not an innocent solvent.
Tollens distinguishes aldehydes from ordinary ketones, but exceptional reducing compounds exist.
Aldol requires α-H; Cannizzaro requires an aldehyde without α-H.
Hofmann degradation shortens the carbon chain by one.
Benzene usually undergoes substitution, not ordinary alkene-like addition.

Self-check¶

Show how Grignard reagents prepare 1°, 2° and 3° alcohols.
Distinguish aldehyde and ketone chemically.
Write equations and conditions for Hofmann, aldol, Cannizzaro and haloform reactions.
Explain Markovnikov orientation using carbocation stability.
Complete a reaction sequence by identifying functional groups and reagents.

Final 60-Minute Revision Plan¶

শেষ ৬০ মিনিটের রিভিশন পরিকল্পনা¶

Time	Task
0–8 min	Atomic formulas, isotopes, Bohr and quantum numbers
8–15 min	Periodic trends and exceptions
15–22 min	Bond comparison, H-bond and Lewis structures
22–29 min	Defect hierarchy, Schottky/Frenkel, p/n-type
29–37 min	`Kc`, `Kp`, Le Chatelier and Hess rules
37–44 min	Zero/first order equations and half-life
44–51 min	pH, `Kw`, M/m/N and dilution
51–60 min	Organic named-reaction decision map

Master formula sheet¶

Topic	Formula
Atomic structure	`A = Z + N`; shell capacity `= 2n²`
Radiation/Bohr	`ΔE = hν`; `mvr = nh/2π`
Periodicity	`Zeff = Z − S`; covalent radius `= ½ bond length`
Equilibrium	`Kp = Kc(RT)^Δn`
Thermochemistry	`ΔH = H(products) − H(reactants)`
Rate	`rate = k[A]^x[B]^y`
Zero order	`[A]t = [A]0 − kt`
First order	`k = (2.303/t)log([A]0/[A]t)`; `t½ = 0.693/k`
Water/pH	`Kw=[H⁺][OH⁻]`; `pH+pOH=14` at 25°C
Concentration	`M=mol/L solution`; `m=mol/kg solvent`; `N=eq/L`
Dilution	`M₁V₁=M₂V₂`

Source coverage note¶

This expanded guide covers the eight source-sheet areas used by Chemistry Sprint: Atomic Structure; Periodic Table; Chemical Bonds; Defects in Solids and Semiconductors; Chemical Equilibrium and Thermochemistry; Chemical Kinetics; pH and Concentration of Solution; and Organic Chemistry. Added explanations are intended to clarify the source, not replace the original PDFs where exact wording or diagrams are required.