Chapter 8 — Study of Compounds

What are elements, compounds and mixtures?
An element is a pure chemical substance made of one kind of atom. Elements cannot be broken down into simpler substances by ordinary chemical means. Examples are iron, oxygen or gold. A compound is formed when two or more different elements chemically combine in fixed, definite proportions. The atoms in a compound are bonded together and the compound shows properties different from those of the constituent elements. For example, sodium is a soft, reactive metal and chlorine is a greenish poisonous gas, but together they form sodium chloride, a white crystalline compound with very different properties.

How mixtures differ from compounds
A mixture is a physical combination of two or more substances in which each substance retains its original chemical identity and properties. The components of a mixture can be present in any proportion and can usually be separated by physical methods such as filtration, decantation, distillation or chromatography. For example, sand mixed with salt can be separated by dissolving the salt in water and filtering the sand. In contrast, to separate a compound into its elements requires a chemical reaction.

Types of mixtures
Mixtures may be homogeneous (uniform composition throughout), such as salt solution or air, or heterogeneous (non-uniform), such as a mixture of iron filings and sulphur powder. Homogeneous mixtures of two or more substances are also called solutions when one substance dissolves in another. Understanding whether a sample is a mixture or a compound tells you what methods to use to separate or analyse it, and what properties to expect.

Properties and identification
Compounds have a fixed composition — a chemical formula — for example water is always H2O. Their physical properties (melting point, boiling point, solubility) and chemical behaviour (reactivity) are specific to that compound. Mixtures, by contrast, show variable compositions and properties depending on their components. To identify whether a substance is a compound or a mixture you can attempt physical separation, check for change in properties on simple mixing, or perform chemical tests — for instance, electrolysis of water yields hydrogen and oxygen only by chemical decomposition, indicating it is a compound.

Why this matters
Distinguishing between these categories is the first step in chemical analysis and practical work. It helps students choose correct laboratory techniques, predict behaviour of materials in everyday life, and appreciate how chemical bonding produces new materials with new properties. This foundation supports later topics like chemical reactions, stoichiometry and materials science.

What is valency?
Valency is a simple way to describe how atoms combine: it represents the number of electrons an atom will lose, gain, or share to reach a stable electron arrangement, usually resembling the nearest noble gas. For many main-group elements this leads to the octet rule: atoms tend to achieve eight electrons in their outermost shell (except hydrogen and helium which aim for two). The valency of an element governs the ratios in which atoms combine and therefore determines the chemical formula of the compound formed.

Electron shells and distribution
Electrons occupy energy levels or shells around the nucleus, labeled K, L, M, N. Each shell has a maximum capacity (K holds 2, L holds 8, M holds 18 but commonly we treat 8 for many simple cases). The outermost shell contains valence electrons which participate in bonding. For example sodium has electron configuration 2,8,1 — one valence electron — so it tends to lose one electron and exhibit valency 1. Oxygen with 2,6 has six valence electrons and tends to gain two to complete an octet, giving valency 2.

Predicting valency
There are practical rules used at this stage: for main group elements valency is often equal to the number of electrons in the outer shell for metals (which lose electrons), and for non-metals it is often 8 minus the number of valence electrons (the number they need to gain to reach eight). Examples: chlorine (2,8,7) needs one electron → valency 1; carbon (2,4) can share four electrons → valency 4. These rules work well for class 10 level compounds and for writing simple formulae.

Electron-dot and Lewis structures
To visualise valence electrons, we use dot structures (Lewis symbols) showing dots for outer electrons around the chemical symbol. These diagrams help show how atoms share or transfer electrons in bonding: Na• and •Cl combine by transfer to make Na+ and Cl−. Lewis structures are helpful to write simple covalent molecules and to deduce lone pairs that affect molecular shape.

Limitations and further ideas
Valency is a useful practical tool but has limits: some elements show variable valency (transition metals) and some compounds do not obey the simple octet. Nevertheless, for most compounds studied in Class 10 the valency concept and electron arrangement provide reliable guidance for predicting formulas, naming compounds and understanding basic bonding types. Practice with many examples will make recognizing patterns quick and accurate.

How ionic bonds form
Ionic bonding takes place when atoms of very different electronegativities interact, typically a metal and a non-metal. The metal atom tends to lose one or more electrons to become a positively charged ion (a cation), while the non-metal gains those electrons to become a negatively charged ion (an anion). This transfer creates opposite charges, and the strong electrostatic attraction between these oppositely charged ions is what holds the compound together. For example, sodium atom (Na) loses one electron to form Na+ and chlorine atom (Cl) gains one to form Cl−; Na+ and Cl− attract each other to give NaCl.

Formation of crystal lattices
Ionic compounds do not exist as separate molecules; instead they form extended three-dimensional crystal lattices where each ion is surrounded by ions of opposite charge. This arrangement maximises attraction and minimises repulsion. The lattice structure explains many macroscopic properties: strong ionic attractions require large amounts of energy to break, hence ionic solids typically have high melting and boiling points and are hard. The repeating pattern also gives characteristic crystal shapes and cleavage planes in minerals.

Physical properties and explanations
Ionic solids are usually brittle: when a force shifts layers, like-charged ions can be forced next to each other, producing strong repulsion that cracks the crystal. In solid state ions are fixed in place, so ionic solids do not conduct electricity. When melted or dissolved in water, ions become mobile and can carry charge; therefore molten ionic compounds and aqueous ionic solutions conduct electricity. Solubility varies: many ionic compounds dissolve readily in polar solvents like water because water molecules stabilise separated ions by solvation or hydration, but some ionic solids have very high lattice energies and are sparingly soluble.

Writing ionic formulas
To write the formula of an ionic compound, balance the total positive and negative charges so the overall charge is zero. This may be done using valencies and the crossover method: write the symbol of each ion with its valency (ignoring sign), then cross the valencies to become subscripts of the opposite ion. Finally simplify by dividing by the greatest common factor if necessary. Example: Al3+ and O2− → Al2O3. Polyatomic ions (like SO42− or NO3−) behave as single charged units and require brackets when more than one is present: Ca2+ + 2NO3− → Ca(NO3)2.

Chemical reactivity and examples
Ionic bonding explains simple displacement reactions: a more reactive metal displaces a less reactive metal from its ionic compound because it can more readily lose electrons. Ionic compounds are common in salts, minerals, and many everyday materials. Practising writing ionic equations and predicting properties from ion charges strengthens understanding of how ionic bonding controls material behaviour.

Nature of covalent bonds
Covalent bonding occurs when two atoms share one or more pairs of electrons so that each atom attains a stable electron arrangement. This type of bonding is common between non-metal atoms. The shared electron pair(s) occupy the region between the two nuclei and hold the atoms together by mutual attraction to the shared electrons. Depending on how many pairs are shared, covalent bonds are classified as single (one pair), double (two pairs) or triple (three pairs).

Bond strength and bond length
Single bonds are generally longer and weaker than double bonds, which in turn are longer and weaker than triple bonds. The greater the number of shared pairs, the stronger the bond and the shorter the bond length. These differences influence the chemical reactivity and physical properties of molecules. For example, nitrogen gas (N≡N) has a strong triple bond and is relatively inert under normal conditions.

Molecules versus network covalent structures
Covalent bonding gives rise to two broad categories: discrete molecular covalent substances and giant covalent (network) solids. Molecular covalent substances consist of separate molecules held together by covalent bonds and weaker intermolecular forces; they often have low melting and boiling points (e.g., H2O molecules, CO2). Giant covalent substances like diamond and graphite form extensive networks of covalent bonds throughout the solid, resulting in very high melting points and distinct properties such as extreme hardness or electrical conductivity (graphite conducts due to delocalised electrons).

Polarity in covalent bonds
Sharing of electrons is not always equal. When atoms of different electronegativity form a bond, the shared electrons are drawn closer to the more electronegative atom, creating a polar covalent bond with partial charges (δ+ and δ−). The extent of polarity affects solubility, boiling point and intermolecular interactions. When polar bonds are arranged asymmetrically in a molecule, the molecule may be polar overall; symmetric arrangements can cancel bond dipoles and produce non-polar molecules.

Lewis structures and resonance
Lewis dot structures represent covalent compounds showing bonds as pairs of dots or lines and lone pairs as dots. Some molecules are best described by resonance: multiple valid Lewis structures where electrons are delocalised, for example in carbonate ion CO32−. Understanding resonance explains stability, equal bond lengths and reactivity patterns.

Applications and examples
Covalent bonding underlies the chemistry of gases, organic molecules, many biomolecules and polymers. Water’s hydrogen bonds (a form of strong intermolecular attraction due to polar covalent O–H bonds) give it unusual properties crucial for life. Learning to draw Lewis structures and predict shapes and polarity builds strong reasoning for later chemistry topics.

Purpose and basic steps
Writing chemical formulae expresses the exact composition of compounds in terms of symbols and numbers. For ionic compounds the goal is to write the simplest whole-number ratio of ions that balances total charge; for covalent molecules the formula lists actual atoms present. The general steps are: identify constituent elements or ions and their valencies or charges, determine the ratio that balances charge, write the formula using subscripts and reduce to the simplest whole-number ratio.

Using valency and crossover method
The crossover method is a practical classroom tool for ionic formulae. Write the symbol of each ion with its valency (ignoring sign), then cross the valencies to become subscripts of the opposite ion. Finally simplify by dividing by the greatest common factor if necessary. Example: Al3+ and O2− → Al2O3 (cross over 3 and 2 to make subscripts 2 and 3). For polyatomic ions, treat the group as a single unit and use brackets when its subscript is more than one, e.g., calcium nitrate: Ca2+ and NO3− → Ca(NO3)2.

Polyatomic ions and their handling
Many common ions contain more than one atom (sulphate SO42−, nitrate NO3−, carbonate CO32−, ammonium NH4+). Learn common polyatomic ions and their charges; they behave like single charged units in formula writing. When multiple polyatomic ions are needed, enclose them in parentheses with the appropriate subscript: for two sulfate ions with aluminium: Al3+ + SO42− → Al2(SO4)3.

Prefixes and covalent naming
For covalent molecular compounds, prefixes indicate the number of each atom: mono-, di-, tri-, tetra- etc. (we often omit mono- for the first element). From names like carbon dioxide (CO2) or dinitrogen pentoxide (N2O5) write formulas directly from the prefixes. Remember that prefix-based names apply to molecular (covalent) compounds, not to ionic compounds.

Common pitfalls and checks
Avoid leaving ionic charges in the final formula; ensure subscripts are reduced to simplest ratio. Check total positive and negative charges balance to zero. For covalent molecules ensure the formula matches the name’s prefixes or the molecular structure given. Practice with many examples including salts with polyatomic ions and transition metal ions with specified oxidation states (e.g., iron(III) chloride FeCl3) to build speed and accuracy.

Why consistent names matter
Chemical nomenclature gives a standard way to name substances so chemists everywhere know what a name means. For Class 10, focus on correct naming of binary ionic compounds, polyatomic salts, simple oxyacids and covalent molecules using common IUPAC-inspired rules that are widely used in school chemistry.

Naming binary ionic compounds
Binary ionic compounds are made of two elements: a metal cation and a non-metal anion. Name the cation (metal) first using its element name and then the anion by taking the root of the non-metal’s name and adding the suffix "-ide". Examples: NaCl sodium chloride; CaO calcium oxide. When the cation is a polyatomic ion, use its name: NH4Cl is ammonium chloride.

Transition metals and oxidation states
Many transition metals have variable oxidation states, so specify the oxidation number in Roman numerals in parentheses after the metal name. For instance FeCl2 is iron(II) chloride because iron has oxidation state +2; FeCl3 is iron(III) chloride for +3. This removes ambiguity and ensures the formula and name correspond uniquely.

Oxyanions and oxyacids
Polyatomic ions containing oxygen (oxyanions) have systematic name patterns: for a pair of ions of the same element with different numbers of oxygen atoms, use the suffixes "-ate" for the more oxygen-rich form and "-ite" for the fewer-oxygen form (e.g., NO3− nitrate, NO2− nitrite). For acids derived from oxyanions: change -ate to -ic and -ite to -ous and add the word "acid" (SO42− sulphate → H2SO4 sulphuric acid; SO32− sulphite → H2SO3 sulphurous acid). Binary acids (hydrogen + non-metal), when in aqueous solution, use the form hydro- + stem + -ic acid: HCl (aq) hydrochloric acid.

Covalent compounds and prefixes
Covalent binaries use Greek prefixes to denote numbers of atoms: mono-, di-, tri-, tetra- etc. Name first element (omit mono- if only one) followed by prefix+stem of second element ending in -ide: CO carbon monoxide; CO2 carbon dioxide; PCl5 phosphorus pentachloride. These names are used for molecular (covalent) substances.

Salts and bases
Bases are typically named as metal hydroxides: NaOH sodium hydroxide, Ca(OH)2 calcium hydroxide. Salts are named cation followed by anion (with its usual name) — e.g., KNO3 potassium nitrate, NH4Cl ammonium chloride. When acid or base names are known, salt names follow directly.

Practice and verification
To avoid mistakes, practise converting names to formulas and formulas to names for a range of compounds including those with polyatomic ions and transition metals. Always check charge balance and oxidation states where relevant. Learning common ion names (nitrate, sulphate, carbonate, hydroxide, ammonium) is especially useful for fast identification and writing of formulas.

Defining acids and bases
At the Class 10 level, an acid can be described as a substance that produces H+ (or H3O+) ions when dissolved in water; a base produces OH- ions or accepts H+ ions. This operational viewpoint explains many observable properties in aqueous solutions.

Properties and classification
Acids generally taste sour (but tasting is unsafe in lab), turn blue litmus red, and may react with metals to produce hydrogen gas. Bases feel soapy or slippery and turn red litmus blue. Strong acids and bases ionise fully in water giving high concentrations of H+ or OH- and thus conduct electricity well; weak acids and bases ionise partially and show smaller conductivity. Salts can be neutral, acidic or basic depending on the strengths of the parent acid and base.

Naming acids and producing salts
Binary acids (H + non-metal) in aqueous solution are named hydro- + stem + -ic acid (for example, HCl is hydrochloric acid). Acids derived from oxyanions follow the -ate → -ic and -ite → -ous rule (for example, NO3- nitrate → HNO3 nitric acid). Neutralisation reactions between acids and bases produce salt and water: HCl + NaOH → NaCl + H2O. The type of salt formed depends on the reacting acid and base: the metal from the base pairs with the anion from the acid.

pH and indicators
pH is a measure of hydrogen ion concentration in solution; it is a practical scale used to compare acidity and basicity. A pH less than 7 indicates acidity, pH greater than 7 indicates basicity, and pH 7 is neutral at standard conditions. Indicators change colour at different pH ranges and are useful for titrations: litmus (red in acid, blue in base), phenolphthalein (colourless in acid, pink in base) and methyl orange (red in acid, yellow in base).

Applications and everyday examples
Neutralisation has many uses: antacid tablets neutralise excess stomach acid; lime (calcium hydroxide) is used to neutralise acidic soils; industrial effluents are treated by neutralisation before release. Understanding acids, bases and salts connects classroom reactions to real-life processes, environmental protection and medicines.

Safety notes
Many acids and bases are corrosive - handle with care, use protective gloves and goggles, and dilute acids by adding acid to water slowly to avoid splashing and heat evolution. Always follow teacher instructions for safe handling and disposal.

What are oxidation numbers?
Oxidation numbers (or states) are theoretical charges assigned to atoms in compounds by a set of rules that help track electron transfer in chemical reactions. They provide a useful bookkeeping tool to identify which atoms gain or lose electrons in a reaction. While oxidation numbers do not always represent real charges on atoms in covalent molecules, they are invaluable for identifying oxidations and reductions and balancing redox reactions.

Rules for assigning oxidation numbers
Memorise practical rules: the oxidation number of an element in its free state is zero (O2, H2, metal elements). For monoatomic ions, oxidation number equals the ion charge (Na+ = +1, Cl- = -1). Hydrogen is usually +1 when bonded to non-metals and -1 in metal hydrides. Oxygen is usually -2 except in peroxides (-1) or when bonded to fluorine. The sum of oxidation numbers in a neutral compound is zero; in a polyatomic ion it equals the ion’s charge. Using these rules, assign numbers to atoms case by case.

Identifying oxidation and reduction
Oxidation is the process of increasing oxidation number (loss of electrons), while reduction is the process of decreasing oxidation number (gain of electrons). A redox reaction includes both processes: one species is oxidised and another is reduced. For example, in Zn + Cu2+ → Zn2+ + Cu, zinc goes from 0 to +2 (oxidised) and copper goes from +2 to 0 (reduced). The substance that causes oxidation (accepts electrons) is an oxidising agent, while the substance that causes reduction (donates electrons) is a reducing agent.

Balancing redox equations (basic idea)
Balancing redox reactions may be done by ensuring conservation of mass and charge, often using half-reaction method at higher levels. For class 10, focus on identifying which atoms change oxidation number and confirming electron transfer balances. Check both atom balance and net charge balance after combining oxidation and reduction changes.

Applications and examples
Redox reactions are central to chemistry and everyday life: rusting of iron (oxidation), combustion, corrosion, and electrochemical cells (batteries) rely on redox processes. In analysis, redox titrations are used to determine concentrations of reactive substances. Practising assignment of oxidation numbers across varied compounds (oxides, halides, sulphates) helps build intuition and prepares you for equation balancing and interpretation of reaction mechanisms.

Concept and definition
Coordinate bonding, also called dative bonding, is a special type of covalent bond in which the pair of electrons shared in the bond originates from the same atom (the donor). The atom that accepts the electron pair (the acceptor) has an empty orbital capable of accommodating the shared pair. Although the origin of the shared pair is different, once formed the coordinate bond behaves like any covalent bond in terms of length, strength and directionality.

How coordinate bonds form
Coordinate bonds commonly occur when a Lewis base (molecule or ion with a lone pair) donates its lone pair to a Lewis acid (an electron-deficient species or a positively charged ion). Examples in inorganic chemistry include the formation of the ammonium ion NH4+ from NH3 and H+: ammonia supplies a lone pair to the proton, forming a N–H bond where both electrons are donated by nitrogen. In complex ions transition metal cations (which have empty d-orbitals) accept electron pairs from ligands such as NH3, H2O or Cl- to form coordination complexes.

Notation and representation
Coordinate bonds are often drawn using an arrow from donor atom to acceptor (for example :NH3 → H+). In many structures the bond is simply drawn as a normal line with an explanation that it is coordinate in origin. In naming coordination complexes, the ligand names and oxidation state of the central metal are important; these topics are introduced later in higher classes but basic examples like [Cu(NH3)4]2+ help illustrate the idea that several donor molecules can bond to a single metal centre.

Properties and importance
Coordinate bonds contribute to the stability and reactivity of many compounds. In biological systems, coordinate bonding is central: haem in haemoglobin binds oxygen to iron through coordinate interactions; enzymes often use metal centres bound by coordinate bonds to activate substrates. In industry, catalysts and complex ions rely on the formation and breaking of coordinate bonds.

Distinguishing features and examples
In NH4+ the additional N–H bond formed by donation of an electron pair from nitrogen makes ammonium positively charged and tetrahedral in shape. In metal complexes ammonia acts as a neutral ligand donating a lone pair to the metal. Carbon monoxide can act as a ligand by donating a lone pair from carbon to metal and accepting back pi-electron density, showing how coordinate bonding can involve complex electron interactions. Understanding coordinate bonding helps explain structures beyond simple ionic or covalent models and connects inorganic chemistry to real-world catalysts and biological chemistry.

VSEPR theory basics
Valence Shell Electron Pair Repulsion (VSEPR) theory is a practical model to predict the shape of molecules. The key idea is simple: electron pairs (bonding pairs and lone pairs) around a central atom repel one another and thus arrange themselves in three-dimensional space to minimise repulsion. The arrangement of electron pairs determines the molecular geometry, and lone pairs exert greater repulsion than bonding pairs, which causes slight distortions in bond angles.

Electron pair domains and common geometries
Start by counting electron pair domains (each bonding pair or lone pair counts as one domain) around the central atom. Typical geometries follow from the number of domains: two domains give a linear arrangement (180°) as in CO2; three domains give trigonal planar (120°) as in BF3; four domains give tetrahedral (109.5°) as in CH4; five domains give trigonal bipyramidal and six give octahedral. When lone pairs are present, the observed molecular shape may differ from the electron-pair geometry: for instance, four domains with one lone pair gives trigonal pyramidal (NH3) rather than tetrahedral, and with two lone pairs it gives bent (H2O).

Effect of lone pairs on bond angles
Lone pairs occupy more space than bonding pairs because they are localised closer to the nucleus and are not shared between atoms. This extra repulsion compresses bond angles between bonding pairs. For example, the ideal tetrahedral angle is 109.5°, but in NH3 (one lone pair) the H–N–H angle is about 107°, and in H2O (two lone pairs) the H–O–H angle is ~104.5°. Recognising how lone pairs change angles is important for predicting exact molecular shapes.

Polarity and shape
Molecular polarity depends on both bond dipoles and molecular geometry. Even if individual bonds are polar, symmetric arrangements can cancel dipoles resulting in non-polar molecules (e.g., CO2). Asymmetric shapes where bond dipoles do not cancel produce polar molecules with permanent dipole moments (e.g., H2O). Understanding shape therefore helps predict physical properties like solubility, boiling point and interactions with electric fields.

How to practice
To determine molecular shape: (1) draw Lewis structure, (2) count electron pair domains on central atom, (3) determine electron-pair geometry, (4) account for lone pairs to find molecular geometry and estimate bond angles. Practice with a variety of examples — linear, trigonal planar, tetrahedral, trigonal pyramidal, bent, trigonal bipyramidal and octahedral — and sketch three-dimensional representations. Visual models or simple ball-and-stick kits help internalise the spatial arrangements for accurate predictions.

Bond polarity and electronegativity
Bond polarity arises when two bonded atoms differ in electronegativity — the tendency of an atom to attract shared electrons. The more electronegative atom pulls the bonding electrons closer, creating a partial negative charge (δ-) while the other atom becomes partially positive (δ+). The result is a bond dipole with a direction from δ+ to δ-. Examples include the O–H bond in water or the C–Cl bond in chloromethane.

Molecular polarity: vector sum of bond dipoles
Molecules may be polar or non-polar depending on the arrangement of individual bond dipoles. Each bond dipole is a vector quantity and the overall molecular dipole is the vector sum. In symmetrical molecules like carbon tetrachloride (CCl4) or carbon dioxide (CO2), bond dipoles cancel out due to the geometry, making the molecule non-polar overall. In asymmetrical molecules like water, bond dipoles add to produce a net dipole, making the molecule polar.

Consequences of polarity
Polarity strongly affects physical properties. Polar molecules interact via dipole-dipole forces and, if they contain H bonded to N, O or F, they can form hydrogen bonds — a particularly strong type of dipole attraction. Hydrogen bonding raises boiling points and influences solubility: polar substances dissolve well in polar solvents like water, while non-polar substances dissolve in non-polar solvents like hexane. For example, water’s polarity and hydrogen bonding explain its relatively high boiling point and its role as a good solvent for many ionic and polar substances.

Predicting molecular polarity
To predict polarity: draw the Lewis structure, determine the molecular geometry using VSEPR, mark bond polarities using electronegativity difference and then perform a vector addition of bond dipoles. If the result is zero, the molecule is non-polar; if not, it is polar. Keep in mind that small differences in geometry or bond strength can influence the final dipole.

Examples and practical tests
Simple tests can indicate polarity: polar liquids align in an electric field and can attract a small stream of water or cause a charged rod to attract or repel them differently than non-polar liquids. Observing solubility trends, boiling point comparisons and surface tension behaviour also gives clues about polarity. Understanding polarity is essential for explaining many chemical and biological phenomena, from dissolution of salts to folding of proteins.

Understanding solubility
Solubility is the ability of a substance (solute) to dissolve in a solvent to form a homogeneous solution. Whether a substance dissolves depends on intermolecular interactions: polar solutes and ionic compounds dissolve more readily in polar solvents like water because solvent molecules stabilise ions or polar molecules through ion-dipole or dipole-dipole interactions. Non-polar solutes dissolve better in non-polar solvents due to London dispersion forces. Temperature and pressure also affect solubility: for most solids solubility in water increases with temperature, while gas solubility typically decreases with rising temperature and increases with pressure.

Properties linked to bonding type
The type of bonding in a compound largely determines its physical properties. Ionic compounds typically have high melting and boiling points because of strong electrostatic forces between ions; they are hard but brittle and conduct electricity when molten or in solution due to mobile ions. Covalent molecular compounds often have low melting and boiling points because intermolecular forces (van der Waals, dipole-dipole, hydrogen bonding) are weaker than ionic attractions; they generally do not conduct electricity. Network covalent solids (diamond, quartz) have very high melting points and are hard because of an extended network of covalent bonds. Metallic bonding gives metals properties like electrical conductivity, malleability and metallic lustre due to delocalised electrons.

Solubility rules and practical use
There are useful solubility patterns useful in laboratory and industry: nitrates (NO3-) and most alkali metal salts are soluble; chlorides are generally soluble except AgCl, PbCl2; sulphates are mostly soluble except BaSO4, CaSO4 and PbSO4. These rules help predict whether a precipitate will form when two solutions are mixed and guide separation techniques. Understanding solubility is important in preparing medicines, purifying compounds and treating wastewater.

Relation to reactivity and applications
Solubility influences reaction rates and outcomes — ionic reactants dissolve to produce free ions which can collide and react; insoluble reactants may react slowly at interfaces. Solubility controls bioavailability of drugs, mobility of pollutants in soil and water, and the behaviour of salts in agriculture. Lattice energy and hydration energy govern whether an ionic compound dissolves: if hydration energy exceeds lattice energy, dissolution is favoured.

Experimental observations and measurement
Solubility curves show how solubility changes with temperature and are used to guide crystallisation for purification. Conductivity measurements indicate the presence of ions in solution. Practical understanding of solubility and properties of compounds enables correct choice of solvents, methods of separation and safe handling in lab and industry.

Hydrides: classification and properties
Hydrides are compounds of hydrogen with other elements and are generally classified by bonding type. Ionic (saline) hydrides form with very electropositive metals (alkali and some alkaline earth metals), e.g., NaH; they contain H- ions and react vigorously with water to release hydrogen gas. Covalent hydrides form between hydrogen and non-metals and include a wide range of organic and inorganic molecules (CH4, NH3). Metallic or interstitial hydrides form when transition metals incorporate hydrogen atoms into their lattice and often exhibit metallic properties rather than discrete ionic or covalent behaviour. The chemical reactivity and physical state of hydrides vary widely: ionic hydrides are typically solids and strong reducing agents; covalent hydrides can be gases, liquids or solids at room temperature.

Oxides: types and reactions
Oxides are compounds composed of oxygen and another element. They are commonly divided into basic oxides (usually metal oxides like Na2O, CaO), acidic oxides (non-metal oxides like CO2, SO2), amphoteric oxides (like Al2O3 and ZnO which react both with acids and bases) and neutral oxides (such as CO and N2O which show little reaction with acids or bases). Basic oxides react with acids to form salts and water; acidic oxides react with bases to form salts and water. Amphoteric oxides can react both ways. These reactions are important in metallurgy, environmental chemistry and industrial processes.

Preparation and uses
Hydrides are useful reducing agents and in hydrogen storage technologies; metal hydrides store hydrogen reversibly in some applications. Oxides are ubiquitous: metal oxides are used as refractories, pigments, ceramics and catalysts; oxides of carbon and sulphur are important pollutants produced by combustion and industrial processes. Controlling oxide formation (corrosion prevention) and converting harmful oxides into less damaging forms are major industrial tasks.

Chemical behaviour and environmental context
Oxides like SO2 and NOx are atmospheric pollutants that can form acid rain by producing sulphuric and nitric acids in the presence of water. Carbon dioxide is a greenhouse gas that contributes to climate change. Understanding the chemical reactivity of oxides informs environmental mitigation strategies, e.g., removal of SO2 from flue gases or catalytic converters that reduce NOx emissions.

Laboratory observations and tests
Some hydrides release H2 with water or acids; ionic hydrides react violently with water while covalent hydrides like CH4 are much less reactive. Oxides may show characteristic reactions: CO2 produces effervescence with limewater, metal oxides often dissolve in acids, and some oxides change colour on heating or reduction. Learning typical reactions and properties helps identify hydrides and oxides in practical work.

Importance of qualitative tests
Qualitative chemical tests are quick ways to identify the presence of particular ions or gases in a sample. These tests rely on characteristic reactions such as colour changes, precipitate formation, gas evolution, flame colours and chemical odours. In the laboratory and in examinations, being able to select and interpret appropriate tests is a practical skill linking theory with observation.

Tests for cations
Several cations give distinctive results: flame tests are simple and useful — sodium gives a bright yellow flame, potassium gives a lilac flame (best seen through cobalt glass to mask sodium), calcium gives brick-red, and copper gives blue-green. Precipitation tests with specific reagents are widely used: adding sodium hydroxide can produce characteristic coloured precipitates (Fe2+ gives green, Fe3+ gives brown, Cu2+ gives blue). Silver(I) ions produce a white precipitate of AgCl with chloride ions, soluble in ammonia; lead(II) gives a white precipitate with chloride that is insoluble in cold water.

Tests for anions
Common anion tests include adding dilute nitric acid followed by silver nitrate for chloride: a white precipitate AgCl confirms chloride. Sulphate ions are identified by acidifying with dilute acid and adding barium chloride; a white precipitate of BaSO4 indicates sulphate. Carbonates give effervescence with dilute acid due to CO2 evolution; the gas turns limewater milky. Sulphites reduce iodide in acid to give iodine, and nitrate can be tested by reduction methods or specific reagent tests.

Gas tests and observations
Hydrogen shows a characteristic 'pop' on ignition (the pop test) due to its combustibility. Oxygen relights a glowing splint due to supporting combustion. Carbon dioxide extinguishes a flame and turns limewater milky. Chlorine gas bleaches litmus paper and has a choking smell; ammonia has a pungent smell and turns moist red litmus blue. Always use small samples and proper safety when testing gases.

Procedure and safety
Perform tests carefully: acidify solutions when required to remove interfering ions (for example carbonate can give false positives for sulphate unless removed). Confirm results using more than one test where possible (for example both precipitation and solubility tests). Use fume hood for toxic gases, keep distance, and never inhale test gases. Dispose of chemical waste as directed and clean glassware after tests.

Interference and confirmation
Some tests give similar results for different ions, so follow systematic test schemes: first test for carbonate (effervescence), then for chloride (AgNO3), then for sulphate (BaCl2 after acidification). Observations such as colour, solubility in reagents (ammonia, dilute acids) and changes on heating help confirm the identity. Practise interpreting real experimental observations to become confident in qualitative analysis.

Common laboratory preparation techniques
Preparing inorganic compounds in the school laboratory uses a handful of reliable methods: direct combination (elements react to form a compound), neutralisation (acid + base → salt + water), precipitation (mixing solutions to form an insoluble salt), thermal decomposition (heating a compound to produce simpler substances) and displacement reactions (a more reactive element displaces a less reactive one from its compound). Understanding which method suits a target compound guides the design of safe, successful experiments.

Crystallisation and purification
Crystallisation is a simple and widely used technique to obtain pure solid salts from their solutions. The general procedure is to prepare a saturated solution, filter to remove insoluble impurities, gently evaporate or cool to encourage crystal formation, then collect and dry the crystals. Recrystallisation — dissolving the solid again in a minimal amount of hot solvent and then allowing it to crystallise — increases purity. Filtration separates solids from liquids, while distillation separates liquids by boiling point differences. These basic operations underpin many laboratory preparations and purifications.

Practical steps and observations
Record all observations carefully: colour changes, formation of precipitates, evolution of gases, temperature changes, and solubility behaviour. Write balanced equations for the reactions to understand stoichiometry and to calculate theoretical yields. For example, preparing copper sulphate by reacting copper oxide with dilute sulphuric acid produces a blue solution which on evaporation yields blue crystals; write the equation CuO + H2SO4 → CuSO4 + H2O and confirm mass changes during crystallisation.

Safety and reagent handling
Use appropriate quantities and concentration of reagents to minimise hazards. Wear goggles, apron and gloves; work in a well-ventilated area or fume hood for volatile or harmful reagents; label containers clearly. When diluting acids, always add acid slowly to water to avoid exothermic splashing. Proper disposal of chemical waste prevents contamination and environmental harm; neutralise acidic or basic waste when required and follow school protocols for hazardous materials.

Troubleshooting and good technique
If crystals do not form, check that the solution was saturated and allowed to cool slowly; slow cooling often gives larger, purer crystals. If precipitates are fine and difficult to filter, use flocculating agents or pressure filtration. Contamination can arise from dirty glassware; wash and rinse thoroughly. Practise careful measurement, slow addition of reagents, and patience — many successful preparations require controlled conditions and attention to detail.

Empirical vs molecular formula
The empirical formula of a compound is the simplest whole-number ratio of atoms of each element in the compound. The molecular formula gives the actual number of atoms of each element in a molecule and may be a multiple of the empirical formula. Determining empirical formula from percentage composition or mass data is an important quantitative skill: convert percentage to mass, mass to moles, then divide by the smallest number of moles to obtain the simplest integer ratio.

Step-by-step empirical formula calculation
Typical steps: assume a 100 g sample so percentages become grams; divide grams by atomic masses to get moles of each element; divide all mole values by the smallest mole value to get a simple ratio; if necessary multiply all numbers by a common factor to remove fractional values. For example, a sample with 40% C, 6.7% H and 53.3% O converts to 3.33 mol C, 6.7 mol H and 3.33 mol O; dividing by 3.33 yields C1H2O1, empirical formula CH2O.

Basic stoichiometry using balanced equations
Stoichiometry relates amounts of reactants and products using mole ratios from balanced chemical equations. Convert given masses to moles using molar masses, use mole ratios to find moles of desired substance, then convert back to mass if required. For gas calculations at the same conditions, volumes are proportional to moles. For Class 10, practice problems usually involve simple integer mole ratios and direct conversions using molar mass = sum of atomic masses.

Practical examples and yields
In the laboratory, compare theoretical yield (calculated from stoichiometry) with actual experimental yield to compute percent yield. Losses can occur due to incomplete reactions, side reactions, or loss during purification. Understanding stoichiometry helps plan reagent quantities and estimate product amounts for reactions such as neutralisation, precipitation and displacement.

Common sources of error and tips
Careful weighing and purity of reagents affect accuracy. In empirical formula problems watch for rounding errors when mole ratios are close to simple fractions (for example 1.5 → multiply all by 2). Keep track of units, use correct atomic masses and show clear working steps. Practise converting masses to moles and applying mole ratios across a variety of problems to build speed and confidence in quantitative reasoning.

General safety rules
Working safely in the chemistry laboratory is essential. Before any experiment read instructions carefully, wear protective equipment (goggles, apron, closed shoes and gloves where required), and remove dangling jewellery or loose clothing. Know the location and proper use of safety equipment such as eyewash stations, safety showers, fire extinguishers and first-aid kits. Never work alone in the laboratory; always have a teacher or supervisor present.

Safe handling of reagents
Label all containers clearly and use small quantities of chemicals in school experiments. Store incompatible chemicals separately (for example oxidisers away from flammable organics, acids away from bases). When diluting strong acids, always add acid slowly to water while stirring to avoid violent heat release and splashing. Use pipettes, spatulas and forceps to handle reagents rather than bare hands, and never taste or directly inhale chemicals. Work in a well-ventilated area or under a fume hood when vapours may be produced.

Glassware and heating
Inspect glassware for cracks before use. Use appropriate supports and clamps for heating apparatus, and heat gently to avoid sudden boiling or bumping. Never point the open end of a test tube towards yourself or others when heating; always direct away. Use tongs or heat-resistant gloves to handle hot equipment. When breaking glass, follow correct disposal procedures — never place broken glass in regular bins.

Waste disposal and spill management
Disposal of chemical waste must follow school guidelines. Neutralise acidic or basic waste when required and segregate non-hazardous and hazardous waste. For spills, alert others, contain the spill if you are trained and use appropriate absorbents; for dangerous spills evacuate and seek teacher help. If a chemical contacts skin or eyes, rinse immediately with plenty of water and seek medical attention.

Behaviour and emergency response
No running or horseplay in the lab. Keep benches tidy and clean up after experiments. In case of fire, alert the teacher, switch off heat sources if safe, and use a fire extinguisher if trained; for clothing on fire use a safety blanket to smother flames. Know emergency contact numbers and the protocol for serious incidents. Being cautious and prepared prevents accidents and ensures a productive learning environment.

Compounds in daily life
Chemical compounds are part of everyday life: water, table salt, sugars, medicines, fertilizers, detergents and plastics are all compounds with distinct properties that determine their uses. Understanding why a substance dissolves, reacts or is stable helps make informed choices: for example, why certain cleaners remove grease (non-polar solvents dissolve non-polar oils) or why some medicines are formulated as salts for better solubility and absorption.

Environmental impacts
Some compounds have harmful environmental effects. Oxides of nitrogen and sulphur produced from combustion react with water in the atmosphere to form acids, causing acid rain that damages crops, buildings and aquatic systems. Persistent organic pollutants, chlorinated solvents and some pesticides resist degradation and bioaccumulate up the food chain, harming wildlife and human health. Fertilizer runoff rich in nitrates and phosphates causes eutrophication in water bodies, creating algal blooms and oxygen depletion.

Chemical solutions and mitigation
Chemistry provides methods to reduce and remediate environmental harm. Flue-gas desulfurisation removes SO2 from industrial emissions, catalytic converters reduce NOx from vehicles, and wastewater treatment neutralises and precipitates harmful ions. Neutralisation with lime (CaCO3) can correct acidic soils and waters. Understanding the chemistry of pollutants allows design of targeted treatments and safer alternatives; for instance, biodegradable detergents reduce persistent pollution compared with older formulations.

Safe use and disposal at home
Household chemicals should be used as directed and disposed of responsibly. Avoid pouring oils and chemicals down drains that feed municipal systems; many localities have hazardous waste collection days. Read labels for hazard warnings and keep incompatible substances separate. Simple chemical knowledge helps prevent accidental mixing of incompatible cleaners that could produce toxic gases (for example bleach with acids produces chlorine gas).

Role of students and future studies
Students who understand compounds and their environmental roles can make better choices and contribute to sustainable practices: reduce single-use plastics, choose eco-friendly products, and support policies that limit emissions. This unit prepares students for advanced study in chemistry, environmental science and related careers, and fosters awareness that chemical knowledge can be applied responsibly to solve real-world problems.