Chapter 5 — Language of Chemistry

Chemistry uses a compact language of names, symbols and formulae to describe substances and their changes. This language helps to write facts clearly and briefly. Instead of writing 'two hydrogen atoms and one oxygen atom make water' every time, we write H2O. Letters like H and O are chemical symbols for elements. Numbers written as subscripts (for example the 2 in H2) tell how many atoms of each element are in a molecule or formula unit. Chemical equations use these symbols and formulae to show reactions: the left side lists reactants and the right side lists products with an arrow between them. A balanced chemical equation also shows conservation of mass because the same number of each atom appears before and after the reaction. Learning the language means learning common element symbols, the rules for writing symbols (first letter capital, second letter small), how to use subscripts and parentheses, and how to read and write short descriptions of reactions. The language also includes names and symbols for polyatomic groups (like SO4 2-) and standard laboratory symbols that warn about hazards. Using this language allows students and scientists anywhere in the world to understand the same reaction or recipe for a compound. Practically, learning these basics enables you to read labels, follow experiments safely, and communicate observations precisely so that others can repeat them. Building fluency comes from practice: converting names to symbols, writing formulae, and expressing reactions in words and then in balanced symbolic form.

An element is a pure substance made of only one kind of atom. Each element has a unique chemical symbol, usually one or two letters taken from its English or Latin name. Symbols must be written with the first letter capitalised and the second letter in small case where applicable: for example, C for carbon, Ca for calcium, and Fe for iron. Many elements are metals (like iron and copper) and others are non-metals (like oxygen and nitrogen). Some symbols come from Latin names — for instance Na for sodium (natrium) and K for potassium (kalium) — so it is important to memorise the common ones used in school chemistry. Elements may exist alone as single atoms (in rare situations) or as molecules of two or more identical atoms; common diatomic elements are H2, O2, N2, Cl2 and F2. Learning a small list of element symbols (for H, O, C, N, Na, K, Ca, Fe, Cl, S, P, Mg) gives the tools to write many chemical names and formulas. Recognising whether an element is a metal or non-metal helps to predict the type of compound it will form: metals tend to form ionic compounds with non-metals, while non-metals often form covalent compounds among themselves. Practice writing element symbols correctly, connecting them with their names and common uses, and noting special cases where Latin names are the source of symbols.

Compounds are substances formed when atoms of different elements join chemically in fixed ratios. A chemical formula is the short way to show which elements are present and how many atoms of each are present. For molecular compounds, the formula tells the number of atoms in a single molecule, for example CH4 shows one carbon and four hydrogen atoms in a methane molecule. For ionic compounds the formula often shows the simplest whole-number ratio of ions in the crystal lattice, for example NaCl shows sodium and chloride in a 1:1 ratio. Subscripts written after an element symbol indicate the number of atoms of that element in one unit; if no subscript is written, the number is one. Parentheses are used when a group of atoms repeats in the formula: Ca(OH)2 indicates two OH groups attached to one calcium atom. It is important to write formulas correctly because small changes change the substance: H2O (water) is different from H2O2 (hydrogen peroxide).

In school exercises you will meet formulas with polyatomic ions — groups of atoms that act as one charged unit such as sulfate (SO4 2-) or nitrate (NO3-). When such a group appears more than once, parentheses enclose the group and the subscript after the parenthesis shows how many groups there are. For example, Al2(SO4)3 has three sulfate groups and two aluminium ions. Another point is reduction to simplest terms: ionic formulas should show the lowest whole-number ratio (for example, empirical ratio), so if you find a formula like Mg2O2 you simplify to MgO. For molecular compounds, the molecular formula already gives the number of atoms per molecule and may not be reducible. Learning to read and write formulas includes practice converting from names to formulas and back, checking that ionic formulas are electrically neutral by balancing positive and negative charges, and placing parentheses correctly for repeating groups. These skills form the foundation for understanding chemical composition, calculating relative masses, and describing reactions in symbolic form.

Matter can be a pure substance (an element or a compound) or a mixture of substances. Pure substances have fixed composition and definite properties: pure water always has formula H2O and specific boiling and freezing points under set conditions. Mixtures contain two or more substances that are not chemically bonded and so their composition may vary. Mixtures are of two types: homogeneous (solutions) and heterogeneous. A homogeneous mixture looks uniform throughout — for example, salt dissolved in water or sugar syrup; you cannot see separate parts. A heterogeneous mixture shows different parts, like sand and iron filings or oil and water; you can often see or separate the components.

Knowing whether a sample is a mixture or a pure substance helps decide how to separate or test it. Physical methods separate mixtures without changing chemical identity: filtration separates insoluble solids from liquids (sand from water), evaporation recovers a dissolved solid by removing the solvent (salt from saltwater), distillation separates liquids with different boiling points (to obtain pure water from a solution), and magnetic separation removes magnetic solids (iron filings) from non-magnetic ones (sand). Decanting can separate immiscible liquids or suspensions by carefully pouring off the top layer. Chromatography (an advanced school technique) separates coloured mixtures into their components. In everyday life, these ideas explain why we use methods like sieving, filtering, or boiling to clean water, prepare food or recycle materials. In the laboratory, selecting an appropriate method depends on the properties of the components (solubility, particle size, boiling point, magnetism). Practice classifying common examples—air, milk, brass, seawater—and choosing a suitable separation technique to develop practical skills and logical thinking about mixtures.

Valency is a basic idea used to predict how atoms join to form compounds. It is the combining capacity of an atom — the number that shows how many bonds it can form or how many electrons it can gain, lose or share to achieve a stable arrangement. For Class 7, learn common valencies: hydrogen (1), oxygen (2), nitrogen (3), carbon (4), chlorine (1), sodium (1) and calcium (2). Valency helps determine simplest ratios of atoms in compounds. For example, when sodium (valency 1) combines with oxygen (valency 2), two sodium atoms are needed to balance one oxygen atom, giving Na2O. For ionic compounds, valency corresponds to the charge on the ion (e.g., Ca2+ has valency 2). A practical method to write formulas from valencies is the criss-cross method: write the symbols with their valencies and cross the numbers down as subscripts of the opposite ion, then simplify to the lowest whole-number ratio. In covalent compounds valency shows how many electrons an atom shares with others, like carbon sharing four electrons to form four bonds in CH4. Remember that valency for many elements at this level is fixed and is a helpful counting tool; later you will study electrons and bonding in more detail. Practice drawing simple combinations and checking that total positive and negative valencies balance for ionic compounds to ensure correct formulas.

Naming compounds correctly links the chemical language to everyday words and helps you convert between names and formulas. For ionic compounds made of a metal and a non-metal, the name is usually formed by naming the metal first and then the non-metal with an '-ide' suffix. For example, NaCl is sodium chloride and CaO is calcium oxide. When a compound contains a polyatomic ion — a group of atoms acting as one charged unit — use the ion's established name rather than adding '-ide': for example, Na2SO4 is sodium sulfate and Ca(NO3)2 is calcium nitrate. Be careful with spelling: sulfate not sulphate in modern notation for many curricula, but learn the version your teacher prefers.

Covalent compounds, formed between two non-metals, use prefixes to indicate the number of atoms of each element present: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), etc. The prefix 'mono-' is usually omitted for the first element, so CO is carbon monoxide (not monocarbon monoxide), CO2 is carbon dioxide, and PCl3 is phosphorus trichloride. For more complex names with multiple atoms of the first element, use the appropriate prefix: N2O5 is dinitrogen pentoxide. Practice both directions: given a name, write the formula; given a formula, write the name. Remember also that ionic names do not use prefixes to show numbers because the ratio is implied by charges, while covalent names use prefixes because molecules may have different numbers of atoms. Learning these naming rules makes it easier to read labels, identify compounds in experiments, and communicate clearly in chemistry tasks at school and at home.

Molecular formula and empirical formula are two useful ways to represent composition. The molecular formula gives the actual number of each type of atom present in one molecule of a substance. For example, ethanol has molecular formula C2H6O meaning two carbons, six hydrogens and one oxygen in each molecule. The empirical formula gives the simplest whole-number ratio of atoms; for ethanol the molecular formula C2H6O cannot be simplified further, but for glucose (C6H12O6) the empirical formula is CH2O after dividing by 6. For hydrogen peroxide (H2O2) the empirical formula is HO because dividing by 2 yields the 1:1 ratio. Finding an empirical formula from percentage composition is a stepwise process used in many school problems: assume a 100 g sample so percent values become grams, convert grams to moles using atomic masses, divide each mole value by the smallest mole value to obtain ratios, and round to the nearest whole number to obtain the simplest integer ratio.

Structural models go beyond formulas by showing how atoms are connected. In Class 7, simple dot-and-cross diagrams and ball-and-stick sketches help visualise covalent bonds. A dot-and-cross diagram represents electrons from different atoms as dots or crosses to show shared pairs in a bond, for example H—H or H:O:H for water (showing two shared pairs). A ball-and-stick model uses labelled circles (balls) for atoms and lines or sticks for bonds, showing spatial arrangement in a simplified way. These models help explain why some molecules have certain shapes or properties, even if the real arrangement is three-dimensional. Practising conversion among molecular formula, empirical formula and simple structural sketches builds students’ ability to read chemical information and relate it to real substances encountered in the laboratory and everyday life.

Chemical reactions describe how substances (reactants) change into new substances (products). We use chemical equations to write reactions: reactants are placed on the left, products on the right, with an arrow between. Each substance is shown by its correct chemical formula. For example, the combustion of methane is written as CH4 + 2O2 → CO2 + 2H2O. This symbolic form is compact and carries detailed information about the types and numbers of atoms involved.

Central to chemical equations is the law of conservation of mass: atoms are neither created nor destroyed, so the count of each type of atom must be equal on both sides of the equation. Balancing an equation means placing whole-number coefficients before formulas so that the numbers of atoms match. Change coefficients only — never change subscripts — because subscripts change the substance. A useful approach is to balance one element at a time, often leaving hydrogen and oxygen for last when they appear in several compounds. Balanced equations can be used later to compare relative amounts of reactants and products, a skill developed further with moles and masses. For Class 7 practice, focus on writing correct formulas, identifying reactants and products, and balancing simple equations for combination, decomposition, displacement, and combustion reactions. Always check your work by counting atoms of each element on both sides. Writing equations clearly, with state symbols (s, l, g, aq) where appropriate, helps communicate experimental observations and shows understanding of how matter changes during reactions.

Classifying reactions helps us understand patterns and predict products. Several simple types are useful for Class 7. Combination (synthesis) reactions occur when two or more simple substances join to form one product: A + B → AB. An example is N2 + 3H2 → 2NH3. Decomposition reactions are the reverse: a single compound breaks into two or more simpler substances, often with heating: AB → A + B (e.g., 2HgO → 2Hg + O2). Single displacement (or replacement) reactions occur when a more reactive element replaces a less reactive one in a compound: A + BC → AC + B, such as Zn + 2HCl → ZnCl2 + H2. Double displacement reactions involve two compounds swapping partners: AB + CD → AD + CB; an example is AgNO3 + NaCl → AgCl (white precipitate) + NaNO3. Combustion reactions involve a substance burning in oxygen with release of heat, commonly producing oxides; burning a hydrocarbon often gives CO2 and H2O (e.g., CH4 + 2O2 → CO2 + 2H2O).

Each type has observable signs: evolution of gas, temperature change (heat produced or absorbed), colour change, formation of a precipitate or a new smell. In many school experiments you can use these signs to identify the type of reaction and write the symbolic equation. Classification also helps in predicting whether a reaction will occur; for instance, a more reactive metal will displace a less reactive metal from its salt. Practise reading short descriptions of experiments, deciding the reaction type, writing balanced equations and explaining the observable changes. This structured approach builds both theoretical understanding and practical laboratory sense for young chemists.

Balancing chemical equations becomes easier with a clear method and regular practice. Start by writing correct formulas for reactants and products. Count atoms of each element on both sides. Pick coefficients to equalise the counts — choose elements that appear in only one reactant and one product first, and leave hydrogen and oxygen for last if they occur in several compounds. If you end up with fractional coefficients, multiply all coefficients by a common factor to obtain whole numbers. Always finish by recounting to confirm balance. Practise with reactions you meet in daily life: combustion of fuels, acid–base reactions, displacement reactions and decomposition reactions. For example, balancing the combustion of methane: CH4 + O2 → CO2 + H2O becomes CH4 + 2O2 → CO2 + 2H2O after balancing.

Applying chemical language to everyday materials builds understanding. Identify common household substances and write their main chemical names or formulas: table salt (NaCl), baking soda (NaHCO3), vinegar's active component acetic acid (CH3COOH), lime (CaO) or limewater (Ca(OH)2). Explain a simple reaction such as the bubbling when vinegar and baking soda mix: NaHCO3 + CH3COOH → CH3COONa + H2O + CO2 — this equation can be balanced and described in words to show gas evolution (CO2) that causes bubbling. Classroom projects often ask students to research an item, write its chemical components and show a simple reaction symbolically. Doing many such exercises develops accuracy in formula writing, equation balancing and translating observations into chemical notation useful for higher classes and everyday science literacy.

Laboratory safety is essential every time you handle chemicals or equipment. Standard safety symbols appear on bottles, containers and laboratory posters so you can quickly recognise hazards and follow correct behaviour. Important symbols include the flame (flammable) which warns that the substance can catch fire easily and must be kept away from open flames, sparks and heat sources. The corrosive symbol (often shown as a test tube pouring liquid on a hand or metal) warns that the substance can burn skin or damage materials; immediate washing and protective gloves/goggles are required. The toxic symbol (skull and crossbones) denotes poisons that must not be ingested or inhaled. The oxidiser symbol warns that the substance can cause or intensify fires when in contact with organic materials. The irritant/harmful symbol indicates milder hazards such as skin or eye irritation. A separate symbol for 'wear eye protection' tells you to wear safety goggles whenever instructed.

Basic safe laboratory practice includes wearing goggles and aprons, tying back long hair, not tasting any chemicals, using a wafting motion to smell gases rather than inhaling directly, labelling all containers clearly, handling hot glassware with tongs, and following instructions for disposing chemical waste. Never mix chemicals unless told to by the teacher and never leave a lit burner unattended. Know the location of safety equipment such as the eyewash station, first aid kit, fire blanket and fire extinguisher. In case of spills, follow classroom procedures: alert the teacher, contain the spill if safe, and clean up using recommended materials. Practising these safety rules protects you and others in the laboratory and builds responsible habits important for science work at school and beyond.