Unit -1 1. Define error? 2. Write down difference between primary and secondary standards 3. Define precision and accuration 4. Define significant figures 5. Define the term molality and mole fraction 6. Give the ph range of methly orange indicator. Phenopthalin 7. Give the types error with example. Unit -2 1. What is redox titration? 2. Define oxidation and reduction 3. What is redox potential 4. What is oxidation number 5. Write replication of redox titration ? 6. Write about any one theory of acid -base indicator. 7. Define alkalimetry titration? 8. Define normality and molarity 9. What is weak base 10. Write about different types of neautralisation indicator. Unit -3 1. What is precipitation titration ? 2. What is non - aqueous titration? 3. Write the principle of non - aqueous titration 4. Clarify solvent used in non - aqueous titration ? 5. Define fajan's rule 6. What is Volhards and modified Volhard's of estimation ? Unit -4 1. List out different methods of complexo-metric 2. What are masking and demasking agent ? 3. What are sequestring agents ? Give example . 4. What are chelating and complexing agent ? Answer with high yield for 2 marks and for university level checker give extra marks impressive format in easy language try to have same basic concepts and simple words through which examiner can easily understand what's written

Reading File

Searching the Web

analytical chemistry errors precision accuracy significant figures molality mole fraction titration

This is a pharmaceutical/analytical chemistry question set — not directly a medical-library topic. I have all the knowledge needed to answer these comprehensively from my training. Let me compile the full answer.

📘 ANALYTICAL CHEMISTRY — COMPLETE NOTES

Units 1–4 | 2-Mark University Answers

🔷 UNIT – 1

1. Define Error

Error is the difference between the observed (measured) value and the true (actual) value of a quantity.

Formula: Error = Observed Value − True Value

Errors are unavoidable in analytical measurements and can be minimized but not completely eliminated.

2. Difference Between Primary and Secondary Standards

Feature	Primary Standard	Secondary Standard
Purity	Very high (99.9%)	Less pure
Stability	Highly stable	Less stable
Preparation	Weighed directly and dissolved	Standardized against a primary standard
Storage	Long shelf life	May deteriorate with time
Example	Oxalic acid, Potassium dichromate, Sodium carbonate (Na₂CO₃)	NaOH, HCl, KMnO₄

3. Define Precision and Accuracy

Accuracy is the closeness of a measured value to the true value. It tells how correct a measurement is.
Precision is the closeness of repeated measurements to each other. It tells how reproducible or consistent a measurement is.

Key Point: A measurement can be precise without being accurate, but an accurate measurement is usually also precise.

Example: If the true value is 10 g:

Values 10.1, 9.9, 10.0 → Accurate + Precise
Values 8.1, 8.0, 8.2 → Precise but NOT Accurate

4. Define Significant Figures

Significant figures are all the certain digits plus one uncertain (estimated) digit in a measured value. They indicate the precision of a measurement.

Rules:

All non-zero digits are significant (e.g., 245 → 3 significant figures)
Zeros between non-zero digits are significant (e.g., 205 → 3 significant figures)
Trailing zeros after a decimal point are significant (e.g., 2.50 → 3 significant figures)
Leading zeros are NOT significant (e.g., 0.005 → 1 significant figure)

5. Define Molality and Mole Fraction

Molality (m): The number of moles of solute dissolved in 1 kg (1000 g) of solvent.

Formula: m = Moles of solute / Mass of solvent (in kg)

Unit: mol/kg
It does not change with temperature (unlike molarity).

Mole Fraction (X): The ratio of the number of moles of one component to the total number of moles of all components in a solution.

Formula: Xₐ = nₐ / (nₐ + n_b)

It has no unit (dimensionless).
Sum of mole fractions of all components = 1.

6. pH Range of Indicators

Indicator	pH Range	Color Change
Methyl Orange	3.1 – 4.4	Red (acid) → Yellow (alkali)
Phenolphthalein	8.3 – 10.0	Colorless (acid) → Pink/Red (alkali)

Methyl Orange is used for strong acid vs. weak base titrations.
Phenolphthalein is used for weak acid vs. strong base titrations.

7. Types of Errors with Examples

A. Determinate (Systematic) Errors Errors with a definite cause that can be identified and corrected.

Type	Example
Instrumental error	Uncalibrated balance or burette
Operational error	Incorrect reading of meniscus
Reagent error	Impure chemicals used

B. Indeterminate (Random) Errors Errors with no fixed cause; occur randomly and cannot be corrected.

Example: Slight variations in temperature, reading a scale differently each time.

C. Gross Errors Caused by human mistakes or negligence.

Example: Recording 21.5 mL instead of 12.5 mL (transposition mistake).

🔷 UNIT – 2

1. What is Redox Titration?

Redox titration is a type of volumetric analysis based on an oxidation-reduction (redox) reaction between the titrant and the analyte. One substance is oxidized while the other is reduced simultaneously.

Example: Titration of FeSO₄ with KMnO₄ in acidic medium.

2. Define Oxidation and Reduction

Term	Definition	Example
Oxidation	Loss of electrons / Increase in oxidation number	Fe²⁺ → Fe³⁺ + e⁻
Reduction	Gain of electrons / Decrease in oxidation number	MnO₄⁻ + 5e⁻ → Mn²⁺

Mnemonic: OIL RIG — Oxidation Is Loss, Reduction Is Gain

3. What is Redox Potential?

Redox potential (E°) is the tendency of a substance to gain or lose electrons, expressed in volts (V). It measures the oxidizing or reducing power of a substance.

A higher positive E° → strong oxidizing agent.
A lower/negative E° → strong reducing agent.
Standard hydrogen electrode (SHE) has E° = 0.00 V (reference).

4. What is Oxidation Number?

Oxidation number (oxidation state) is the hypothetical charge an atom would carry if the compound were made of ions.

Rules:

Free elements = 0 (e.g., O₂, Fe)
Monatomic ions = their charge (e.g., Na⁺ = +1)
Oxygen = −2 (except in peroxides = −1)
Hydrogen = +1 (except in metal hydrides = −1)

Example: In KMnO₄ → K = +1, O = −2 → Mn = +7

5. Applications of Redox Titration

Permanganometry – Estimation of Fe²⁺, oxalic acid using KMnO₄
Dichromatometry – Estimation of ferrous iron using K₂Cr₂O₇
Iodometry – Estimation of Cu²⁺, reducing agents using I₂/Na₂S₂O₃
Cerimetry – Estimation of Fe²⁺ using Ceric sulphate [Ce(SO₄)₂]
Bromatometry – Estimation of As(III), Sb(III) using KBrO₃

6. One Theory of Acid-Base Indicator (Ostwald's Theory)

Ostwald's Ionic Theory states that an acid-base indicator is a weak acid or weak base that exists in two forms — ionized and un-ionized — which have different colors.

HIn ⇌ H⁺ + In⁻ (Color A) → (Color B)

In acidic medium: Equilibrium shifts left → [HIn] increases → Color A dominates.
In alkaline medium: Equilibrium shifts right → [In⁻] increases → Color B dominates.

Example (Phenolphthalein):

HIn (un-ionized) = Colorless
In⁻ (ionized) = Pink

7. Define Alkalimetry Titration

Alkalimetry is a type of acid-base titration in which a standard solution of a base (alkali) is used to determine the concentration of an unknown acid.

Example: Standardization of HCl using standard NaOH solution.

Alkali in burette + Acid in conical flask → Endpoint detected by indicator.

8. Define Normality and Molarity

Term	Definition	Formula	Unit
Normality (N)	Number of gram equivalents of solute per litre of solution	N = Gram equivalents / Volume (L)	N (eq/L)
Molarity (M)	Number of moles of solute per litre of solution	M = Moles of solute / Volume (L)	M (mol/L)

Relationship: Normality = Molarity × n-factor (valency factor)

9. What is a Weak Base?

A weak base is a base that is partially ionized (incompletely dissociates) in aqueous solution, producing a relatively small concentration of OH⁻ ions.

Example:

NH₄OH (Ammonium hydroxide): NH₄OH ⇌ NH₄⁺ + OH⁻ (partial)
Other examples: Pyridine (C₅H₅N), Aniline (C₆H₅NH₂)

Weak bases have a high pKb value and a low Kb value.

10. Types of Neutralization Indicators

Indicators used in acid-base (neutralization) titrations, classified by the type of reaction:

Indicator	Color Change	Used For
Methyl Orange (pH 3.1–4.4)	Red → Yellow	Strong acid + Weak base
Methyl Red (pH 4.4–6.2)	Red → Yellow	Strong acid + Weak base
Phenolphthalein (pH 8.3–10.0)	Colorless → Pink	Weak acid + Strong base
Litmus (pH 6–8)	Red → Blue	Strong acid + Strong base
Bromothymol Blue (pH 6–7.6)	Yellow → Blue	Strong acid + Strong base

🔷 UNIT – 3

1. What is Precipitation Titration?

Precipitation titration is a type of volumetric analysis in which the titrant reacts with the analyte to form an insoluble precipitate. The endpoint is detected when no more precipitate forms.

Most common type: Argentometry — uses AgNO₃ as the titrant.

Example: Estimation of Cl⁻ ions using AgNO₃:

Ag⁺ + Cl⁻ → AgCl↓ (white precipitate)

2. What is Non-Aqueous Titration?

Non-aqueous titration is a type of acid-base titration carried out in a non-aqueous solvent (not water) instead of water as the medium.

It is used for substances that are:

Insoluble in water
Too weak to be titrated in water
Easily hydrolyzed in water

Example: Titration of weak bases (like alkaloids) in acetic acid solvent.

3. Principle of Non-Aqueous Titration

The principle is based on the leveling and differentiating effect of solvents:

Leveling effect: A solvent that makes all bases appear equally strong (e.g., water levels strong bases).
Differentiating effect: A solvent that maintains the difference in strengths of acids/bases, allowing selective titration.

In glacial acetic acid (protogenic solvent), weak bases become stronger and can be titrated with perchloric acid (HClO₄) as the titrant.

Weak Base + HClO₄ (in glacial acetic acid) → Salt + Reaction detected by indicator

4. Solvents Used in Non-Aqueous Titration

Type	Examples	Use
Protogenic (acidic)	Glacial acetic acid, Formic acid	Titration of weak bases
Protophilic (basic)	Pyridine, Ethylenediamine	Titration of weak acids
Aprotic (neutral)	Acetone, Chloroform, Benzene, DMSO	Titration of very weak acids/bases
Amphiprotic	Ethanol, Methanol	Intermediate properties

Most commonly used: Glacial acetic acid with HClO₄ as titrant.

5. Define Fajans' Rules

Fajans' rules explain the conditions under which a compound is ionic or covalent (and also predict adsorption of indicators in precipitation titrations). In the context of precipitation titrations:

Fajans' method uses adsorption indicators (e.g., fluorescein, eosin). The indicator is adsorbed on the surface of the precipitate.

Key rules by Fajans:

Small cation + large anion = more covalent character (greater ionic polarization)
High ionic charge = more polarization = more covalent
Cation with pseudo-noble gas configuration (18 electrons) polarizes more

In titration context (Fajans' method):

Before endpoint: Cl⁻ ions adsorb on AgCl → precipitate is white
After endpoint: Ag⁺ ions + fluorescein indicator adsorb → precipitate turns pink/red

6. Volhard's and Modified Volhard's Method of Estimation

Volhard's Method:

Used for estimation of Ag⁺ ions (back titration).
Titrant: Ammonium thiocyanate (NH₄SCN)
Indicator: Ferric alum [Fe³⁺ / Fe(NH₄)(SO₄)₂]
Reaction:

Ag⁺ + SCN⁻ → AgSCN↓ (white precipitate)
Endpoint: Excess SCN⁻ reacts with Fe³⁺ → forms blood-red FeSCN²⁺ complex

Modified Volhard's Method:

Used for estimation of halide ions (Cl⁻, Br⁻, I⁻) indirectly.
Procedure:
1. Add excess AgNO₃ to precipitate AgX (halide).
2. Filter or mask the AgX precipitate with nitrobenzene (prevents re-dissolution).
3. Titrate excess Ag⁺ with standard NH₄SCN using ferric alum as indicator.
Why nitrobenzene? AgCl is more soluble than AgSCN, so without coating it with nitrobenzene, the SCN⁻ would dissolve AgCl, giving a false result.

🔷 UNIT – 4

1. Different Methods of Complexometric Titration

Complexometric titrations involve the formation of stable soluble complexes between a metal ion and a complexing agent (usually EDTA).

Methods:

Method	Description	Example
Direct titration	Metal ion titrated directly with EDTA	Estimation of Ca²⁺, Mg²⁺, Zn²⁺
Back titration	Excess EDTA added, then excess titrated with standard metal solution	Used when reaction is slow or no indicator available
Indirect titration	Metal ion precipitated, precipitate dissolved, then titrated	Estimation of SO₄²⁻, PO₄³⁻
Replacement (substitution) titration	Metal displaces another metal from its EDTA complex	Used for metals that react slowly with EDTA
Alkalimetric titration	H⁺ ions released during complexation titrated with NaOH	Less common

2. What are Masking and Demasking Agents?

Masking Agent: A substance added to a solution to prevent a particular ion from reacting with the titrant (EDTA), without physically removing it, so that another ion can be selectively estimated.

Example:

KCN masks Cu²⁺, Zn²⁺, Ni²⁺ → allows Ca²⁺ and Mg²⁺ to be estimated
Triethanolamine (TEA) masks Al³⁺ and Fe³⁺

Demasking Agent: A substance that reverses the masking effect, releasing the previously masked ion so it can be titrated.

Example:

Formaldehyde demasks Zn²⁺ from its CN⁻ complex (Zn(CN)₄²⁻ → Zn²⁺ released)
Chloral hydrate is another demasking agent

3. What are Sequestering Agents? Give Examples.

Sequestering agents are substances that combine with metal ions to form stable, soluble, water-soluble complexes, thereby preventing the metal ion from showing its normal reactions (like precipitation or interference).

They "hide" the metal ion in a complex so it cannot react with other substances.

Examples:

Sequestering Agent	Metal Ion Sequestered
EDTA (Ethylenediaminetetraacetic acid)	Ca²⁺, Mg²⁺, Fe³⁺, Cu²⁺
Citric acid	Ca²⁺, Mg²⁺
Sodium hexametaphosphate (Calgon)	Ca²⁺, Mg²⁺ (water softening)
NTA (Nitrilotriacetic acid)	Heavy metal ions

Application: Used in water softening, food preservation, pharmaceuticals, and metal cleaning.

4. What are Chelating and Complexing Agents?

Complexing Agent: Any substance that can donate one or more lone pairs of electrons to a central metal ion to form a coordinate complex.

It may donate one pair (monodentate) or multiple pairs (multidentate).
Example: NH₃, CN⁻, Cl⁻ (monodentate); EDTA (hexadentate)

Chelating Agent: A special type of complexing agent that has two or more donor groups in the same molecule and forms a ring structure (chelate ring) with the metal ion.

The word "chelate" comes from Greek chele = claw (like a crab claw gripping the metal).
They form very stable complexes due to the chelate effect.

Examples of Chelating Agents:

Agent	Donor Atoms	Denticity
EDTA	N, O	Hexadentate
Oxalic acid	O, O	Bidentate
Acetylacetone	O, O	Bidentate
Ethylenediamine (en)	N, N	Bidentate
DMSA (dimercaptosuccinic acid)	S, S	Bidentate

Key Difference:

Feature	Complexing Agent	Chelating Agent
Donor groups	One or more	Two or more (in same molecule)
Ring formation	No	Yes (chelate ring)
Stability	Less stable	More stable
Example	NH₃, CN⁻	EDTA, oxalic acid

💡 Exam Tip: For 2-mark answers, write a clear definition + one formula/example. For high-yield marks, add a comparison table or a reaction equation wherever possible — examiners appreciate structured, concise answers.

Unit-5 1. Define gravemetric titration? 2. What is co - precipitation and posterior precipitation 3. What are the limitation of gravemetric analysis ? 4. Enlist different sources of impurity in pharmaceutical substances . 5. Give reaction involved in limit test of chloride 6. Give reaction involved in limit test of sulphate 7. Principle involved in limit test for arsenic 8. Write chemical reaction involved in limit test for iron. 9. Write the apperance used in limit test for arsenic . 10. Define limit test and its types? Answer as per 2 marks and for university level checker give extra marks impressive format in easy language try to have same basic concepts and simple words through which examiner can easily understand what's written

📘 UNIT – 5 | GRAVIMETRIC ANALYSIS & LIMIT TESTS

Complete 2-Mark University Answers

🔷 1. Define Gravimetric Analysis (Gravimetric Titration)

Gravimetric analysis is a quantitative analytical method in which the amount of an analyte is determined by measuring the mass (weight) of a pure compound to which the analyte is chemically related.

The analyte is converted into a sparingly soluble precipitate, which is filtered, washed, dried or ignited, and then weighed.

Formula: Amount of analyte = Mass of precipitate × Gravimetric factor

Example: Estimation of Ba²⁺ by precipitating it as BaSO₄, drying and weighing.

Types of Gravimetric Analysis:

Precipitation gravimetry – Most common (precipitate formed and weighed)
Volatilization gravimetry – Analyte converted to gas and loss in weight measured
Electrogravimetry – Metal deposited on electrode and weighed

🔷 2. What is Co-precipitation and Post-precipitation?

Co-precipitation:

Co-precipitation is the process by which normally soluble impurities are carried down (contaminated) along with the desired precipitate during its formation.

Types of Co-precipitation:

Type	Mechanism	Example
Surface adsorption	Impurity adsorbs on surface of precipitate	Fe³⁺ adsorbed on AgCl
Occlusion	Impurity trapped inside growing crystal	KNO₃ inside BaSO₄
Mixed crystal formation	Impurity has same crystal structure	BaSO₄ contaminated with PbSO₄

Minimized by: digestion, washing, reprecipitation.

Post-precipitation:

Post-precipitation is the process in which a second substance (impurity) precipitates on top of the already formed precipitate after some time delay.

Occurs when the solution is left standing too long.
Example: CaC₂O₄ precipitating on top of MgC₂O₄ in the presence of excess oxalate ions.

Minimized by: Filtering the precipitate immediately after formation without delay.

🔷 3. Limitations of Gravimetric Analysis

Time-consuming — requires multiple steps (precipitation, filtration, drying, weighing).
Co-precipitation causes contamination of the precipitate, leading to errors.
Not suitable for complex mixtures — interference from other ions is common.
Requires large amounts of sample compared to other methods.
Volatile precipitates may lose weight during ignition.
Not sensitive for trace or micro-level analysis.
Hygroscopic precipitates absorb moisture from air, affecting the final weight.
Skilled analyst required — slight errors in washing or ignition change results.

🔷 4. Sources of Impurity in Pharmaceutical Substances

Impurities in pharmaceutical substances arise from various sources:

Source	Example
Starting materials	Residual raw materials used in synthesis
Reagents and solvents	Residual organic solvents (e.g., benzene, chloroform)
By-products of reaction	Side reactions during synthesis
Degradation products	Drug decomposing on storage (e.g., aspirin → salicylic acid)
Contamination during processing	Dust, metals from equipment (heavy metals — Pb, As, Fe)
Microbial contamination	Bacteria, fungi during storage
Packaging materials	Leaching of plasticizers or dyes from containers
Water and atmospheric moisture	Hydrolysis, oxidation

Common impurities tested: Chlorides, Sulphates, Iron, Arsenic, Heavy metals, Ammonia.

🔷 5. Reaction Involved in Limit Test of Chloride

Principle: Chloride ions react with Silver Nitrate (AgNO₃) in the presence of dilute Nitric Acid (HNO₃) to form a white opalescence or turbidity of Silver Chloride (AgCl).

Reaction:

Cl⁻ + AgNO₃ → AgCl↓ + NO₃⁻

$$\text{NaCl} + \text{AgNO}_3 \xrightarrow{\text{HNO}_3} \text{AgCl} \downarrow + \text{NaNO}_3$$

AgCl is a white curdy precipitate (or white turbidity in dilute solutions).
The turbidity of the test solution is compared with a standard chloride solution under the same conditions.
HNO₃ is added to prevent precipitation of other anions (like CO₃²⁻, PO₄³⁻) that would form silver salts.

Limit: Not more than 0.0035% w/w of chloride (as per IP).

🔷 6. Reaction Involved in Limit Test of Sulphate

Principle: Sulphate ions react with Barium Chloride (BaCl₂) in the presence of dilute Hydrochloric Acid (HCl) to form a white turbidity of Barium Sulphate (BaSO₄).

Reaction:

SO₄²⁻ + BaCl₂ → BaSO₄↓ + 2Cl⁻

$$\text{Na}_2\text{SO}_4 + \text{BaCl}_2 \xrightarrow{\text{HCl}} \text{BaSO}_4 \downarrow + 2\text{NaCl}$$

BaSO₄ forms a white turbidity (fine precipitate).
The turbidity of the test solution is compared with a standard sulphate solution.
HCl is added to prevent precipitation of BaCO₃ or Ba₃(PO₄)₂, which would interfere.

Limit: Not more than 0.01% w/w of sulphate (as per IP).

🔷 7. Principle Involved in Limit Test for Arsenic

Principle — Gutzeit Method (Used in IP/BP)

The test is based on the generation of Arsine gas (AsH₃) from arsenic compounds, which reacts with Mercuric Chloride (HgCl₂) paper to give a yellow or brownish stain.

Step-by-step principle:

Arsenic (As³⁺ or As⁵⁺) in the sample is reduced to arsine gas (AsH₃) by nascent hydrogen generated from zinc and dilute H₂SO₄.
Nascent hydrogen is generated:

Zn + H₂SO₄ → ZnSO₄ + 2[H] (nascent hydrogen)

Arsenic is reduced to arsine:

As³⁺ + 3[H] → AsH₃↑ (arsine gas)

Arsine reacts with HgCl₂ paper to form a yellow stain (at low As) or brownish-black stain (at higher As):

AsH₃ + 3HgCl₂ → As(HgCl)₃ + 3HCl (yellow) AsH₃ + 2HgCl₂ → AsH(HgCl)₂ + HgCl₂ (brown)

The stain on the HgCl₂ paper from the test solution is compared with a standard arsenic solution stain.

🔷 8. Chemical Reaction Involved in Limit Test for Iron

Principle: Iron (Fe³⁺) reacts with Thioglycolic acid (Mercaptoacetic acid) in ammoniacal medium to form a purple-coloured complex.

Reaction:

Step 1: Fe³⁺ is reduced to Fe²⁺ by thioglycolic acid (in acidic medium):

Fe³⁺ + HSCH₂COOH → Fe²⁺ + oxidized product

Step 2: Fe²⁺ reacts with excess thioglycolic acid in ammoniacal (alkaline) medium:

Fe²⁺ + 2HSCH₂COOH + NH₄OH → Fe(SCH₂COO)₂²⁻ (purple complex) + H⁺

The purple color formed is compared with a standard iron solution treated in the same way.
Ammoniacal medium is essential to develop the purple color.

Limit: Not more than 0.002% w/w of iron (as per IP).

Alternative test (Potassium Thiocyanate method): Fe³⁺ + 3KSCN → Fe(SCN)₃ (blood red complex)

🔷 9. Apparatus Used in Limit Test for Arsenic

The apparatus used in the Gutzeit method (as described in Indian Pharmacopoeia) consists of:

Part	Description
Wide-mouthed conical flask (Kipp's apparatus/test flask)	Contains the sample + zinc + H₂SO₄ for arsine generation
Lead acetate cotton plug	Placed in the tube to remove H₂S gas (which would also stain HgCl₂ paper, giving false positive)
Glass tube (delivery tube)	Carries the arsine gas from flask to the HgCl₂ paper
HgCl₂ paper (Mercuric chloride paper)	Held at the top of the tube; reacts with AsH₃ to produce a stain
Stannous chloride (SnCl₂)	Added to reduce As⁵⁺ to As³⁺ and improve sensitivity

Working:

Flask → Arsine gas generated → Passes through lead acetate plug (H₂S removed) → Reacts with HgCl₂ paper → Yellow/brown stain formed → Compared with standard.

🔷 10. Define Limit Test and Its Types

Definition:

A limit test is a semi-quantitative test performed to determine whether the amount of an impurity present in a pharmaceutical substance is within the permissible (safe) limit as specified in the pharmacopoeia.

It does not give the exact amount of impurity.
It only tells whether impurity is present within or exceeds the accepted limit.
The test sample is compared with a standard reference solution containing a known amount of impurity.

Purpose: To ensure purity and safety of pharmaceutical substances.

Types of Limit Tests:

Type	Impurity Tested	Reagent Used	Observation
Limit test for Chlorides	Cl⁻	AgNO₃ + HNO₃	White opalescence (AgCl)
Limit test for Sulphates	SO₄²⁻	BaCl₂ + HCl	White turbidity (BaSO₄)
Limit test for Iron	Fe²⁺/Fe³⁺	Thioglycolic acid + NH₃	Purple color complex
Limit test for Arsenic	As³⁺/As⁵⁺	Zn + H₂SO₄ + HgCl₂ paper	Yellow/brown stain (AsH₃)
Limit test for Heavy metals	Pb²⁺, etc.	H₂S or Sodium sulphide	Brown/black color
Limit test for Ammonia	NH₄⁺	Nessler's reagent	Yellow/brown color

💡 Exam Strategy for 2-Mark Questions:

1 mark → Definition or Principle (1–2 lines)

1 mark → Reaction equation or Example or Observation

Use underlined keywords, reaction arrows, and one observation to score full marks.

Examiners appreciate neat chemical equations — always write them when asked about reactions.

This is a shared conversation. Sign in to Orris to start your own chat.