Acids, Bases, and Salts: The Ultimate Guide for IGCSE & O Level Chemistry Mastery
Your complete resource for IGCSE Chemistry 0620, O Level Chemistry 5070, AQA, and Edexcel syllabuses. Packed with definitions, reactions, notes, and exam secrets.
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Part 1: Core Concepts & Definitions
Understanding the fundamental definitions is the first step to conquering this topic. We'll explore the key theories you need to know for your O Level Chemistry 5070 or IGCSE Chemistry 0620 exam[citation:1][citation:6].
1. The Arrhenius Theory (The Foundation)
Proposed by Svante Arrhenius in 1884, this is the simplest definition and the one most frequently tested[citation:1].
- Arrhenius Acid: A substance that increases the concentration of H⁺ ions (protons) when dissolved in water[citation:1]. Example: HCl → H⁺ + Cl⁻
- Arrhenius Base: A substance that increases the concentration of OH⁻ ions (hydroxide ions) when dissolved in water[citation:1]. Example: NaOH → Na⁺ + OH⁻
Key Fact: Free H⁺ ions don't really exist in water. They instantly bond with a water molecule to form the hydronium ion, H₃O⁺[citation:1]. So, HCl(aq) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq).
2. Physical Properties: How to Identify Them
You can often identify acids and bases by their observable properties[citation:4][citation:10].
| Property | Acids | Bases |
|---|---|---|
| Taste | Sour (e.g., lemon juice, vinegar)[citation:5] | Bitter (e.g., soap, baking soda)[citation:5] |
| Touch/Texture | Usually feel like water. | Feel soapy and slippery[citation:4]. |
| Effect on Litmus | Turn blue litmus paper red[citation:5]. | Turn red litmus paper blue[citation:5]. |
| pH Value | pH < 7[citation:4] | pH > 7[citation:4] |
| Reaction with Metals | React with many metals to produce a salt and hydrogen gas (H₂)[citation:2][citation:10]. | Typically do not react with most metals[citation:4]. |
3. What is a Salt?
A salt is an ionic compound formed when an acid reacts with a base, a process called neutralization[citation:5][citation:6].
General Equation: Acid + Base → Salt + Water[citation:2]
For example: Hydrochloric Acid + Sodium Hydroxide → Sodium Chloride + Water[citation:2]
HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)
Salts are made up of positive ions (cations, often from the base) and negative ions (anions, from the acid). Not all salts are neutral—their pH in water depends on the strength of the parent acid and base[citation:3][citation:9].
Part 2: Common Acids & Bases in Daily Life and the Lab
Chemistry isn't just in textbooks. Here are examples you encounter every day[citation:5][citation:8].
🛒 Everyday Acids
- Citric Acid: In lemons, oranges, and other citrus fruits.
- Acetic Acid (CH₃COOH): The main component of vinegar[citation:8].
- Lactic Acid: Found in yogurt and sour milk.
- Tartaric Acid: In tamarind and grapes.
- Carbonic Acid (H₂CO₃): In carbonated soft drinks.
- Hydrochloric Acid (HCl): A strong acid found in our stomach gastric juices[citation:8].
🧼 Everyday Bases
- Sodium Hydroxide (NaOH, Caustic Soda): In drain and oven cleaners[citation:8].
- Calcium Hydroxide (Ca(OH)₂, Slaked Lime): Used in mortar and whitewash[citation:8].
- Ammonia (NH₃): In many household glass and surface cleaners[citation:8].
- Magnesium Hydroxide (Mg(OH)₂): The active ingredient in "Milk of Magnesia," an antacid.
- Sodium Bicarbonate (NaHCO₃, Baking Soda): Used in baking and as a mild cleaning agent.
Part 3: Key Reactions You MUST Know (with Equations)
This is the heart of the exam. Memorize the patterns, not just individual equations[citation:2][citation:10].
🎯 Exam Tip:
For O Level Chemistry 5070 past papers and IGCSE Chemistry 0620 past papers, you are almost always required to write balanced chemical equations. Practice these general patterns!
1. Acid + Metal → Salt + Hydrogen Gas
A classic test for acids[citation:2][citation:10].
- Example: Zinc + Hydrochloric Acid → Zinc Chloride + Hydrogen[citation:10]
Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g) - Why it matters: The "pop" test for hydrogen gas is a frequent O Level Chemistry 5070 ATP (Alternative to Practical) question.
2. Acid + Base (Neutralization) → Salt + Water
The most fundamental acid-base reaction[citation:2][citation:6].
- Example: Sulfuric Acid + Sodium Hydroxide → Sodium Sulfate + Water[citation:8]
H₂SO₄(aq) + 2NaOH(aq) → Na₂SO₄(aq) + 2H₂O(l) - Net Ionic Equation: The core of any strong acid-strong base neutralization is: H⁺(aq) + OH⁻(aq) → H₂O(l)[citation:2].
3. Acid + Metal Carbonate → Salt + Water + Carbon Dioxide
Another key test for acids[citation:4][citation:10].
- Example: Calcium Carbonate (Marble) + Hydrochloric Acid → Calcium Chloride + Water + CO₂
CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g) - Why it matters: The effervescence (fizzing) due to CO₂ is a standard identification test.
4. Acidic/Amphoteric Oxides & Basic Oxides
Oxides also show acid-base character[citation:10].
- Acidic Oxide + Base → Salt + Water: e.g., CO₂ + Ca(OH)₂ (limewater) → CaCO₃ + H₂O[citation:10].
- Basic Oxide + Acid → Salt + Water: e.g., CuO + H₂SO₄ → CuSO₄ + H₂O[citation:10].
- Amphoteric Oxides (e.g., Al₂O₃, ZnO) react with both acids and strong bases[citation:4].
Watch a Past Paper Solved by an Expert
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Part 4: pH, Salt Hydrolysis, and Predicting Salt Acidity
Why is NaCl neutral but NH₄Cl acidic? The answer lies in hydrolysis[citation:3].
The pH of Salts: The 4 Rules
When a salt dissolves, its ions can react with water. This is hydrolysis. Use this table to predict the pH of a salt solution[citation:3]:
| Salt Made From... | Example | pH of Solution | Reason (Hydrolysis) |
|---|---|---|---|
| Strong Acid + Strong Base | NaCl (HCl + NaOH) | = 7 (Neutral) | Neither ion reacts with water. Spectator ions only[citation:3]. |
| Strong Acid + Weak Base | NH₄Cl (HCl + NH₃) | < 7 (Acidic) | The cation (NH₄⁺) from the weak base donates H⁺ to water[citation:3]. |
| Weak Acid + Strong Base | CH₃COONa (CH₃COOH + NaOH) | > 7 (Basic/Alkaline) | The anion (CH₃COO⁻) from the weak acid accepts H⁺ from water, leaving OH⁻[citation:3]. |
| Weak Acid + Weak Base | CH₃COONH₄ | ≈ 7 (Near Neutral)* | Depends on relative strengths (Ka vs Kb). Usually close to 7[citation:3]. |
*Requires comparison of acid and base dissociation constants (Ka, Kb) for exact pH.
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📝 Top 5 Exam Strategies for Acids, Bases, Salts
- Memorize the Reaction Patterns: Acid+Metal, Acid+Carbonate, Neutralization. Write them out daily.
- Know Your Tests: H₂ (pop test), CO₂ (limewater turns milky), NH₃ (turns damp red litmus blue).
- Predict Salt pH: Use the 4 rules from Part 4. This is a high-mark discriminator question.
- Practice Ionic Equations: Especially the net ionic equation for neutralization (H⁺ + OH⁻ → H₂O).
- Use the O Level Chemistry Formula Sheet: Know what's on it and how to use it fast.
📚 Your Ultimate Chemistry Download Hub
Download comprehensive, exam-focused notes for every major topic. Click the links to get the PDFs from Google Drive.
Pro Tip: Combine these with the Save My Exams Notes for IGCSE Chemistry 0620 for unbeatable revision.
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Live Class Schedule: Weekly classes on Stoichiometry & Moles, Organic Chemistry, and Acids & Bases Revision. New batch starting soon. Check the updated syllabus & schedule here.
Watch a full O Level Paper 4 solved: See Prof. Janjowa's problem-solving approach in action.
Part 6: Last-Minute Revision & Cheat Sheets
Crunch time? These resources are your lifesaver.
- Last Minute Revision Notes for O Level Chemistry 5070
- O Level Chemistry Notes PDF (All Chapters)
- How to Calculate Empirical Formulas Easily (Related to salt formulas)
Ultimate Cation/Anion & Gas Test Cheat Sheet
Memorize this table for identification questions in Paper 4 (ATP).
Learn the Periodic Table in a Song!
Knowing Group I & II metals (which react with acids) and Group VII (halogens, part of acids) is crucial. This makes it fun!
Frequently Asked Questions (FAQs)
How can I tell if an unknown substance is an acid, base, or salt?
Step 1: Litmus Test. Blue to Red = Acid. Red to Blue = Base. No change could be a neutral salt (but not always).
Step 2: Reaction Tests. If it's an acid, it will fizz with a carbonate or a reactive metal. A base will feel soapy.
Step 3: pH Measurement. Use universal indicator or a pH probe for a definitive number[citation:4][citation:5].
Is NaCl (table salt) an acid or a base?
Neither. NaCl is a neutral salt. It is formed from a strong acid (HCl) and a strong base (NaOH). When dissolved, neither Na⁺ nor Cl⁻ ions react with water, so the pH remains 7[citation:3][citation:9].
Why does my stomach hurt if I have too much acid, and how do antacids work?
Stomach pain can be from excess hydrochloric acid (HCl). Antacids like magnesium hydroxide (Mg(OH)₂) are bases. They undergo a neutralization reaction with the stomach acid: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O. This reduces acidity and relieves pain[citation:8][citation:10].
What's the difference between a strong acid and a concentrated acid?
This is a common confusion! Strength refers to how completely an acid dissociates into ions. A strong acid (like HCl) dissociates 100%. A weak acid (like CH₃COOH) does not. Concentration is about how much acid is dissolved in a given volume of water. You can have a dilute strong acid or a concentrated weak acid[citation:1][citation:10].
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![Fundamental Concepts & States of Matter • Atom: The smallest particle of an element that can exist, made of a nucleus (protons and neutrons) and electrons orbiting it. • Element: A pure substance consisting of only one type of atom, which cannot be broken down into simpler substances by chemical means. • Compound: A substance formed when two or more different elements are chemically bonded together in a fixed ratio. • Mixture: A substance containing two or more elements or compounds not chemically bonded together. Can be separated by physical means. • Molecule: A group of two or more atoms held together by chemical bonds. • Proton: A subatomic particle found in the nucleus with a relative mass of 1 and a charge of +1. • Neutron: A subatomic particle found in the nucleus with a relative mass of 1 and no charge (0). • Electron: A subatomic particle orbiting the nucleus with a negligible relative mass and a charge of -1. • Atomic Number (Z): The number of protons in the nucleus of an atom. Defines the element. • Mass Number (A): The total number of protons and neutrons in the nucleus of an atom. • Isotopes: Atoms of the same element (same atomic number) but with different mass numbers due to a different number of neutrons. • Relative Atomic Mass ($A_r$): The weighted average mass of an atom of an element compared to $1/12$th the mass of a carbon-12 atom. • Relative Molecular Mass ($M_r$): The sum of the relative atomic masses of all atoms in one molecule of a compound. • Relative Formula Mass ($M_r$): The sum of the relative atomic masses of all atoms in the formula unit of an ionic compound. • Mole: The amount of substance that contains $6.02 \times 10^{23}$ particles (Avogadro's number). • Molar Mass: The mass of one mole of a substance, expressed in g/mol. Numerically equal to $A_r$ or $M_r$. • Empirical Formula: The simplest whole number ratio of atoms of each element in a compound. • Molecular Formula: The actual number of atoms of each element in a molecule. • Solid: Particles are closely packed in a fixed, regular arrangement, vibrate about fixed positions. Definite shape and volume. • Liquid: Particles are closely packed but randomly arranged, can slide past each other. Definite volume, no definite shape. • Gas: Particles are far apart and arranged randomly, move rapidly and randomly. No definite shape or volume. • Melting Point: The specific temperature at which a solid changes into a liquid at a given pressure. • Boiling Point: The specific temperature at which a liquid changes into a gas (vaporizes) at a given pressure. • Sublimation: The direct change of state from solid to gas without passing through the liquid phase (e.g., solid $\text{CO}_2$). • Diffusion: The net movement of particles from a region of higher concentration to a region of lower concentration, due to random motion. • Osmosis: The net movement of water molecules across a partially permeable membrane from a region of higher water potential to a region of lower water potential. 2. Structure & Bonding • Ionic Bond: The electrostatic force of attraction between oppositely charged ions, formed by the transfer of electrons from a metal to a non-metal. • Covalent Bond: A strong electrostatic force of attraction between a shared pair of electrons and the nuclei of the bonded atoms, typically between two non-metals. • Metallic Bond: The electrostatic force of attraction between positive metal ions and delocalised electrons. • Ion: An atom or group of atoms that has gained or lost electrons, resulting in a net electrical charge. • Cation: A positively charged ion (lost electrons). • Anion: A negatively charged ion (gained electrons). • Octet Rule: Atoms tend to gain, lose, or share electrons in order to achieve a full outer electron shell, typically with eight electrons. • Giant Ionic Lattice: A regular, repeating 3D arrangement of oppositely charged ions, held together by strong electrostatic forces. • Simple Molecular Structure: Molecules held together by strong covalent bonds, but with weak intermolecular forces between molecules. • Giant Covalent Structure (Macromolecular): A large structure where all atoms are held together by strong covalent bonds in a continuous network (e.g., diamond, silicon dioxide). • Allotropes: Different structural forms of the same element in the same physical state (e.g., diamond and graphite are allotropes of carbon). • Electronegativity: The power of an atom to attract the electron pair in a covalent bond to itself. • Polar Covalent Bond: A covalent bond in which electrons are shared unequally due to a difference in electronegativity between the bonded atoms. • Hydrogen Bond: A strong type of intermolecular force that occurs between molecules containing hydrogen bonded to a highly electronegative atom (N, O, F). • Van der Waals' forces: Weak intermolecular forces of attraction between all molecules, arising from temporary dipoles. 3. Stoichiometry & Chemical Calculations • Stoichiometry: The study of quantitative relationships between reactants and products in chemical reactions. • Limiting Reactant: The reactant that is completely consumed in a chemical reaction, determining the maximum amount of product that can be formed. • Excess Reactant: The reactant present in a greater amount than required to react with the limiting reactant. • Yield: The amount of product obtained from a chemical reaction. • Theoretical Yield: The maximum amount of product that can be formed from a given amount of reactants, calculated using stoichiometry. • Actual Yield: The amount of product actually obtained from a chemical reaction, usually less than the theoretical yield. • Percentage Yield: $($Actual Yield $/$ Theoretical Yield$) \times 100\%$. • Concentration: The amount of solute dissolved in a given volume of solvent or solution. Often expressed in mol/dm$^3$ (molarity) or g/dm$^3$. • Solute: The substance that dissolves in a solvent to form a solution. • Solvent: The substance in which a solute dissolves to form a solution. • Solution: A homogeneous mixture formed when a solute dissolves in a solvent. 4. Chemical Reactions & Energetics • Chemical Reaction: A process that involves rearrangement of the atomic structure of substances, resulting in the formation of new substances. • Reactants: The starting substances in a chemical reaction. • Products: The substances formed as a result of a chemical reaction. • Word Equation: An equation that uses the names of the reactants and products. • Symbol Equation: An equation that uses chemical symbols and formulae to represent reactants and products, and is balanced. • Balancing Equation: Ensuring the number of atoms of each element is the same on both sides of a chemical equation. • Redox Reaction: A reaction involving both reduction and oxidation. • Oxidation: Loss of electrons, gain of oxygen, or loss of hydrogen. Increase in oxidation state. • Reduction: Gain of electrons, loss of oxygen, or gain of hydrogen. Decrease in oxidation state. • Oxidising Agent: A substance that causes oxidation by accepting electrons (and is itself reduced). • Reducing Agent: A substance that causes reduction by donating electrons (and is itself oxidised). • Exothermic Reaction: A reaction that releases energy to the surroundings, usually as heat, causing the temperature of the surroundings to rise. $\Delta H$ is negative. • Endothermic Reaction: A reaction that absorbs energy from the surroundings, usually as heat, causing the temperature of the surroundings to fall. $\Delta H$ is positive. • Activation Energy ($E_a$): The minimum amount of energy required for reactants to collide effectively and initiate a chemical reaction. • Catalyst: A substance that increases the rate of a chemical reaction without being chemically changed itself, by providing an alternative reaction pathway with a lower activation energy. • Enthalpy Change ($\Delta H$): The heat energy change measured at constant pressure for a reaction. • Standard Enthalpy of Formation ($\Delta H_f^\circ$): The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states under standard conditions. • Standard Enthalpy of Combustion ($\Delta H_c^\circ$): The enthalpy change when one mole of a substance is completely combusted in oxygen under standard conditions. • Hess's Law: The total enthalpy change for a reaction is independent of the route taken, provided the initial and final conditions are the same. 5. Rates of Reaction & Equilibrium • Rate of Reaction: The change in concentration of a reactant or product per unit time. • Collision Theory: For a reaction to occur, reactant particles must collide with sufficient energy (activation energy) and correct orientation. • Factors Affecting Rate: Concentration, pressure (for gases), surface area, temperature, and presence of a catalyst. • Reversible Reaction: A reaction where products can react to reform the original reactants, indicated by $\rightleftharpoons$. • Chemical Equilibrium: A state in a reversible reaction where the rate of the forward reaction is equal to the rate of the reverse reaction, and the concentrations of reactants and products remain constant. • Le Chatelier's Principle: If a change in conditions (temperature, pressure, concentration) is applied to a system at equilibrium, the system will shift in a direction that counteracts the change. 6. Acids, Bases & Salts • Acid: A substance that produces hydrogen ions ($H^+$) when dissolved in water (Arrhenius definition) or a proton donor (Brønsted-Lowry definition). • Base: A substance that produces hydroxide ions ($OH^-$) when dissolved in water (Arrhenius definition) or a proton acceptor (Brønsted-Lowry definition). • Alkali: A soluble base that dissolves in water to produce hydroxide ions ($OH^-$). • Salt: A compound formed when the hydrogen ion of an acid is replaced by a metal ion or an ammonium ion. • Neutralisation: The reaction between an acid and a base (or alkali) to form a salt and water. $H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$. • pH: A measure of the acidity or alkalinity of a solution, defined as $-\log_{10}[H^+]$. Scale from 0 to 14. • Strong Acid: An acid that fully dissociates (ionizes) in water (e.g., HCl, $H_2SO_4$). • Weak Acid: An acid that partially dissociates (ionizes) in water (e.g., $CH_3COOH$). • Strong Base: A base that fully dissociates in water (e.g., NaOH, KOH). • Weak Base: A base that partially dissociates in water (e.g., $NH_3$). • Amphoteric: A substance that can act as both an acid and a base (e.g., aluminium oxide, water). • Titration: A quantitative chemical analysis method used to determine the unknown concentration of a reactant using a known concentration of another reactant. • Indicator: A substance that changes colour over a specific pH range, used to detect the endpoint of a titration. 7. Electrochemistry • Electrolysis: The decomposition of an ionic compound using electrical energy. Requires molten or aqueous electrolyte. • Electrolyte: An ionic compound (molten or dissolved in a solvent) that conducts electricity due to the movement of ions. • Electrodes: Conductors (usually metal or graphite) through which electricity enters and leaves the electrolyte. • Anode: The positive electrode, where oxidation occurs (anions are attracted). • Cathode: The negative electrode, where reduction occurs (cations are attracted). • Faraday's Laws of Electrolysis: Relate the amount of substance produced at an electrode to the quantity of electricity passed through the electrolyte. • Galvanic (Voltaic) Cell: An electrochemical cell that generates electrical energy from spontaneous redox reactions. • Standard Electrode Potential ($E^\circ$): The potential difference of a half-cell compared to a standard hydrogen electrode under standard conditions (1 M concentration, 1 atm pressure for gases, 298 K). • Electrochemical Series: A list of elements arranged in order of their standard electrode potentials, indicating their relative reactivity as oxidising or reducing agents. 8. The Periodic Table • Periodic Table: An arrangement of elements in order of increasing atomic number, showing periodic trends in properties. • Group: A vertical column in the periodic table, containing elements with the same number of valence electrons and similar chemical properties. • Period: A horizontal row in the periodic table, containing elements with the same number of electron shells. • Valence Electrons: Electrons in the outermost shell of an atom, involved in chemical bonding. • Alkali Metals (Group 1): Highly reactive metals, readily lose one electron to form $+1$ ions. React vigorously with water. • Alkaline Earth Metals (Group 2): Reactive metals, readily lose two electrons to form $+2$ ions. • Halogens (Group 17/7): Highly reactive non-metals, readily gain one electron to form $-1$ ions. Exist as diatomic molecules. • Noble Gases (Group 18/0): Unreactive elements with a full outer electron shell, existing as monatomic gases. • Transition Metals: Elements in the d-block of the periodic table, characterised by variable oxidation states, coloured compounds, and catalytic activity. • Metallic Character: Tendency of an element to lose electrons and form positive ions. Increases down a group, decreases across a period. • Non-metallic Character: Tendency of an element to gain electrons and form negative ions. Decreases down a group, increases across a period. • Ionisation Energy: The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous $1+$ ions. • Electron Affinity: The energy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous $1-$ ions. 9. Organic Chemistry • Organic Chemistry: The study of carbon compounds, excluding carbonates, carbides, and oxides of carbon. • Hydrocarbon: A compound containing only carbon and hydrogen atoms. • Saturated Hydrocarbon: A hydrocarbon containing only single carbon-carbon bonds (e.g., alkanes). • Unsaturated Hydrocarbon: A hydrocarbon containing one or more carbon-carbon double or triple bonds (e.g., alkenes, alkynes). • Homologous Series: A series of organic compounds with the same general formula, similar chemical properties, and showing a gradual change in physical properties. • Functional Group: A specific group of atoms within a molecule that is responsible for the characteristic chemical reactions of that molecule. • Alkane: Saturated hydrocarbons with the general formula $C_nH_{2n+2}$. Contain only single bonds. • Alkene: Unsaturated hydrocarbons with the general formula $C_nH_{2n}$. Contain at least one carbon-carbon double bond. • Alkyne: Unsaturated hydrocarbons with the general formula $C_nH_{2n-2}$. Contain at least one carbon-carbon triple bond. • Alcohol: Organic compounds containing the hydroxyl functional group ($-OH$). General formula $C_nH_{2n+1}OH$. • Carboxylic Acid: Organic compounds containing the carboxyl functional group ($-COOH$). • Ester: Organic compounds formed from the reaction of a carboxylic acid and an alcohol, containing the ester linkage ($-COO-$). • Isomers: Compounds with the same molecular formula but different structural formulae. • Structural Isomers: Isomers that differ in the arrangement of their atoms or bonds. • Addition Reaction: A reaction where an unsaturated molecule adds across a double or triple bond, forming a single product. • Substitution Reaction: A reaction where an atom or group of atoms in a molecule is replaced by another atom or group of atoms. • Polymerisation: The process of joining many small monomer molecules together to form a large polymer molecule. • Monomer: A small molecule that can be joined together to form a polymer. • Polymer: A large molecule (macromolecule) formed from many repeating monomer units. • Addition Polymerisation: Polymerisation where monomers add to one another in such a way that the polymer contains all the atoms of the monomer. Usually involves unsaturated monomers. • Condensation Polymerisation: Polymerisation where monomers join together with the elimination of a small molecule (e.g., water). • Cracking: The process of breaking down long-chain hydrocarbons into shorter, more useful hydrocarbons using heat and/or a catalyst. • Fermentation: The anaerobic respiration of yeast, converting glucose into ethanol and carbon dioxide. 10. Analytical Chemistry • Qualitative Analysis: The identification of the components of a sample. • Quantitative Analysis: The determination of the amount or concentration of a component in a sample. • Chromatography: A separation technique based on differential partitioning between a stationary phase and a mobile phase. • Retention Factor ($R_f$): In paper/thin-layer chromatography, the ratio of the distance travelled by the spot to the distance travelled by the solvent front. • Spectroscopy: The study of the interaction of electromagnetic radiation with matter. • Infrared (IR) Spectroscopy: Used to identify functional groups in organic molecules based on their absorption of IR radiation. • Mass Spectrometry: Used to determine the relative molecular mass of a compound and its fragmentation pattern to deduce structure. • Flame Test: A qualitative test for the presence of certain metal ions, which produce characteristic colours when heated in a flame.](https://cambridgeclassroom.com/wp-content/uploads/2024/03/White-And-Purple-Modern-Online-Graphic-Design-Courses-Instagram-Post-4.png)
