Results for "Chlorine"

C2 - Covalent Bonding 1: Bonding in Hydrogen, Chlorine and Hydrogen chloride

Chapters 1 and 2 Study Guide 1.1 Classifying matter (what it is and how we name it) Physical state Solid: fixed shape and volume (particles tightly packed; strong intermolecular forces). Liquid: fixed volume but shape adapts to container (particles close but can move past each other). Gas: no fixed shape or volume — fills container (particles far apart; weak interactions). Pure substance vs mixture Pure substance: one type of “stuff” with fixed composition. Element: pure substance that cannot be chemically broken into simpler substances (e.g., O₂ gas consists of oxygen atoms). Compound: pure substance made of two or more elements chemically bonded in fixed ratio (e.g., H₂O). Mixture: two or more substances physically combined; components keep their chemical identities. Homogeneous mixture: uniform throughout (solution like salt water). Heterogeneous mixture: non-uniform (e.g., sand in water). Particles Atom: smallest unit of an element that retains element identity (e.g., one carbon atom). Molecule: two or more atoms bonded together (could be same element, e.g., O₂, or different, e.g., CO₂). Ion: atom or group of atoms with net electric charge (cation = positive, anion = negative). How to classify — short checklist Is composition fixed? yes → pure substance (element or compound). no → mixture. Is it made of single type of atom? yes → element. If different elements bonded → compound. Is it uniform? yes → homogeneous; no → heterogeneous. 1.2 SI units & common prefixes Base SI units you need: Length: meter (m) Mass: kilogram (kg) — note: kg is base unit (not g), but we often use grams (g). Time: second (s) Temperature: kelvin (K) Amount of substance: mole (mol) Common prefixes (multipliers): kilo- (k) = 10³ = 1,000 e.g., 1 km = 1,000 m centi- (c) = 10⁻² = 0.01 e.g., 1 cm = 0.01 m milli- (m) = 10⁻³ = 0.001 e.g., 1 mm = 0.001 m Tip: memorize that kilo = 1000, centi = 1/100, milli = 1/1000. 1.3 Scientific notation Why: makes very large or very small numbers easier and reduces error. Format: a \times 10^n where 1 ≤ |a| < 10 and n is integer. Examples 0.000345 = 3.45 \times 10^{-4}. 6,200,000 = 6.2 \times 10^{6}. To convert: move decimal left for positive exponent, right for negative. 1.4 Significant figures (sig figs) & uncertainty What sig figs mean: digits in a measurement that are known reliably plus one uncertain digit (last digit). Rules for counting sig figs Nonzero digits are always significant: 245 → 3 sig figs. Zeros between nonzero digits are significant: 2005 → 4 sig figs. Leading zeros (to left of first nonzero) are NOT significant: 0.0072 → 2 sig figs. Trailing zeros in a number with a decimal are significant: 2.300 → 4 sig figs. Trailing zeros in a whole number without a decimal are ambiguous — use scientific notation to show significance: 1200 (ambiguous) → write 1.200×10³ to show 4 sig figs or 1.2×10³ to show 2. Relationship to uncertainty: the last sig fig is the estimated digit — it indicates the measurement’s uncertainty level. 1.5 Sig figs in calculations Multiplication / Division: answer has same number of sig figs as factor with fewest sig figs. Example: 2.5 \times 3.42 = 8.55 → 2 sig figs → round to 8.6. Addition / Subtraction: align decimal places; answer has decimal places equal to the quantity with fewest decimal places. Example: 12.11 + 0.3 = 12.41 → fewest decimals = 1 decimal → round to 12.4. Always follow these rules and only round at the end of multi-step calculations (keeping extra guard digits during intermediate steps). 1.6 Accuracy vs precision Accuracy: how close a measurement is to the true/accepted value. Precision: how repeatable measurements are (how close they are to each other). Illustration (dartboard): All darts close to bullseye → accurate and precise. All darts clustered but far from bullseye → precise but not accurate. Darts spread out but centered on bullseye on average → accurate but not precise. 1.7 Derived units: volume & density Volume: for regular shapes use geometry (e.g., m³ or L = dm³). Common lab units: mL (1 mL = 1 cm³). Density: \rho = \dfrac{\text{mass}}{\text{volume}}. Common units: g/mL or g/cm³ for solids/liquids; kg/m³ in SI. Rearrangements \text{mass} = \text{density} \times \text{volume} \text{volume} = \dfrac{\text{mass}}{\text{density}} Example (lab / irregular object): Object mass = 12.43 g. Volume by displacement = 3.10 mL. Density = 12.43 g ÷ 3.10 mL = 4.0097 g/mL → with sig figs: three sig figs (3.10 has 3 sig figs; 12.43 has 4) → 4.01 g/mL. 1.8 Dimensional analysis (unit conversions) Key idea: multiply by conversion factors that equal 1 (units cancel). Worked example: convert 2.50 miles → meters. 1 mile = 1609.34 m. 2.50\ \text{mi}\times \frac{1609.34\ \text{m}}{1\ \text{mi}} = 4023.35\ \text{m}. Sig figs: 2.50 has 3 sig figs → answer must have 3 sig figs → 4.02 × 10³ m. 1.9 Practice problems (with steps + sig figs) Problem A — Scientific notation Convert 0.000462 to scientific notation. Move decimal 4 places right: 4.62 \times 10^{-4}. Problem B — Multiplication with sig figs Compute (3.60 \times 2.1). Raw: 3.60\times2.1 = 7.56. Sig figs: 3.60 (3 sig figs), 2.1 (2 sig figs) → result 2 sig figs → 7.6. Problem C — Addition with sig figs Compute 12.11 + 0.3 + 0.042. Align decimals; fewest decimal places = 1 (from 0.3) → round final to 1 decimal place. Sum = 12.11 + 0.3 + 0.042 = 12.452 → round to 12.5. Problem D — Density (irregular object) Mass = 24.68 g; initial water in graduated cylinder = 15.0 mL; final = 17.35 mL. Volume displaced = 17.35 − 15.0 = 2.35 mL (note: 15.0 has 3 sig figs so difference has 3 sig figs). Density = 24.68 ÷ 2.35 = 10.500 ≈ 4 sig figs? But check sig figs: mass 24.68 (4 sig figs), volume 2.35 (3 sig figs) → result to 3 sig figs → 10.5 g/mL. Chapter 2 — Atomic structure, periodic table, bonding, naming 2.1 Historic experiments — what they showed and why they matter J.J. Thomson (late 1800s) — cathode ray experiment What he did: passed a beam (cathode ray) through electric & magnetic fields and measured deflection. Observation: beam deflected toward positive plate → beam composed of negatively charged particles. Discovery: existence of the electron — a very small, negatively charged particle present in atoms. Model implication: atoms are not indivisible; they contain subatomic particles. Thomson proposed the “plum pudding” model: a positive “soup” with embedded electrons. Significance: first discovery of subatomic particle; proved atoms have internal structure. Ernest Rutherford (early 1900s) — gold foil experiment What he did: fired alpha particles (positively charged) at very thin gold foil and detected scattering angles. Observation: most alpha particles passed straight through, but a small fraction deflected at large angles; some bounced back. Conclusion: atom is mostly empty space with a tiny, dense, positively charged nucleus that contains most mass — electrons orbit around that nucleus. Model implication: replaced plum pudding with nuclear model (nucleus + orbiting electrons). Why this matters: Rutherford explained the large-angle deflections that Thomson’s model couldn’t; introduced the nucleus concept — foundation for modern atomic structure. 2.2 Atomic number, mass number, isotopes Atomic number (Z): number of protons in nucleus → defines the element. Mass number (A): total number of protons + neutrons in nucleus (integer). Isotopes: atoms of same element (same Z) with different numbers of neutrons (different A). Example: carbon-12 (^12C) and carbon-14 (^14C). Isotopic notation: {}^{A}{Z}X^{\text{charge}} Example: an ion of chlorine with 17 protons and 18 neutrons and −1 charge → {}^{35}{17}\text{Cl}^{-} (35 = 17+18). Average atomic mass: weighted average of isotopic masses using natural abundances (found on periodic table). The periodic table lists average atomic mass (not integer mass numbers) because natural samples are mixtures of isotopes. 2.3 Example: average atomic mass (Neon) Given isotopes (typical values): ^20Ne mass = 19.992440 amu, abundance = 90.48% (0.9048) ^21Ne mass = 20.993847 amu, abundance = 0.27% (0.0027) ^22Ne mass = 21.991386 amu, abundance = 9.25% (0.0925) Average atomic mass: (19.992440)(0.9048) + (20.993847)(0.0027) + (21.991386)(0.0925) = 20.1800\ \text{amu (approx.)} This is the kind of number you’ll see on the periodic table: 20.180 amu. Procedure (general): multiply each isotope mass × its fractional abundance, then sum. 2.4 Writing atomic/ionic symbols and counting particles Given: protons, neutrons, electrons, and charge — decide isotope notation or ion symbol. Example 1: 17 protons, 18 neutrons, 17 electrons → neutral chlorine atom ^35Cl (since A = 17 + 18 = 35). Symbol: {}^{35}_{17}\text{Cl}. Example 2: 11 protons, 12 neutrons, 10 electrons → net charge +1 (lost 1 electron) → sodium ion {}^{23}_{11}\text{Na}^{+} (A = 23). How to check: protons = atomic number (Z) → identifies element. electrons = protons − charge (if charge positive, fewer electrons). mass number A = protons + neutrons. 2.5 Ionic vs covalent (molecular) compounds Ionic compounds Formed when electrons are transferred from a metal to a nonmetal (forming cations and anions). Bonding characterized by electrostatic attraction between oppositely charged ions. Usually full formula is a formula unit (empirical ratio). Often solids with high melting points and conduct electricity when molten or dissolved. Example: NaCl (Na⁺ and Cl⁻). Covalent (molecular) compounds Formed when two nonmetals share electrons to achieve noble gas configuration. Bonds are electron-sharing; molecules have discrete units. Example: H₂O, CO₂. Rule of thumb: metal + nonmetal → usually ionic. nonmetal + nonmetal → usually covalent. 2.6 Using the periodic table to get information From an element’s location: Atomic number → number of protons (top of box). Atomic mass (average) → usually decimal number beneath symbol. Group (column) number → similar chemical behavior and valence electrons. Period (row) → number of electron shells occupied. Metals/Nonmetals/Metalloids: left side metals, right side nonmetals; staircase demarcates metalloids. Common ion charges: Groups 1A → +1, 2A → +2, 7A (halogens) → −1, 6A → −2, etc. Transition metals: central block (d-block). Halogens: Group 17 (F, Cl, Br, I…). Noble gases: Group 18 (He, Ne, Ar…). 2.7 Writing chemical formulas Ionic compounds (simple) Balance total positive and negative charges to get neutral compound. Example: Al³⁺ and O²⁻ → least common multiple of 3 and 2 = 6 → need 2 Al³⁺ (2×+3=+6) and 3 O²⁻ (3×−2=−6) → formula Al₂O₃. With polyatomic ions Treat polyatomic ion as a unit; balance charges. Use parentheses when more than one polyatomic unit is needed: e.g., calcium nitrate = Ca²⁺ + NO₃⁻ → need two NO₃⁻ → Ca(NO₃)₂. Naming Ionic: cation name (metal) first, then anion name (nonmetal with −ide ending) or polyatomic ion name. For transition metals that can have multiple charges, use Roman numeral for charge (iron(III) chloride = FeCl₃). Molecular (binary nonmetal compounds): use prefixes (mono-, di-, tri-, etc.) to show number of each atom (CO₂ = carbon dioxide). Common polyatomic ions (memorize these) NH₄⁺ ammonium NO₃⁻ nitrate SO₄²⁻ sulfate CO₃²⁻ carbonate OH⁻ hydroxide PO₄³⁻ phosphate ClO₄⁻ perchlorate 2.8 Hydrates Definition: a compound that contains water molecules in its crystalline structure: written as \text{salt} \cdot x\text{H}_2\text{O}. Example: copper(II) sulfate pentahydrate = CuSO₄·5H₂O. Naming: name ionic compound then add prefix for water number + “hydrate” (e.g., decahydrate, pentahydrate). Finding empirical hydrate formula (lab procedure) Mass of hydrate (before heating) — measured. Heat to remove water → mass of anhydrous salt measured. Mass of water lost = mass hydrate − mass anhydrous. Convert both masses to moles: moles anhydrous = mass anhydrous ÷ molar mass of anhydrous salt. moles water = mass water ÷ 18.015 g/mol. Compute mole ratio: moles water ÷ moles anhydrous → round to nearest small whole number → that’s x in salt·x H₂O. Worked lab example (complete): (This matches labs you described.) Given: Mass hydrate = 2.564 g Mass anhydrous = 1.622 g Assume the anhydrous formula is CuSO₄ (molar mass = 159.609 g/mol) Steps: Mass water = 2.564 − 1.622 = 0.942 g. Moles anhydrous CuSO₄ = 1.622 g ÷ 159.609 g/mol = 0.010162 mol. Moles water = 0.942 g ÷ 18.015 g/mol = 0.052290 mol. Mole ratio water : salt = 0.052290 ÷ 0.010162 = 5.15 ≈ 5 → formula CuSO₄·5H₂O. Why rounding to whole number: water molecules must be whole; experimental values near whole numbers are rounded to the nearest integer (if close enough — e.g., 2.99 → 3). 2.9 Stoichiometry & dimensional analysis reminders Always write units and let them cancel. Carry at least one extra guard digit through calculations; only round final answer to correct sig figs. For atomic/molecular calculations you’ll often use Avogadro’s number: 6.022\times10^{23} particles/mol. Calculations and practice problems from your list — solved Identifying sig figs — quick answers 0.004500 → 4 significant figures (4500 with leading zeros not significant; trailing zeros after decimal are significant). 1200 → ambiguous; writing as 1.200×10³ shows 4 sig figs; 1.2×10³ shows 2. Example: Putting numbers into/out of scientific notation 7,890,000 → 7.89 \times 10^{6} (3 sig figs if original had 3 sig figs). 3.40\times10^{-5} → 0.0000340. Dimensional analysis multistep example (CH1 style) Problem: Convert 45.0 km/h to m/s. 1 km = 1000 m, 1 h = 3600 s. 45.0\ \text{km/h} \times \dfrac{1000\ \text{m}}{1\ \text{km}} \times \dfrac{1\ \text{h}}{3600\ \text{s}} = \dfrac{45.0\times1000}{3600}\ \text{m/s} = 12.5\ \text{m/s} (sig figs: 45.0 has 3 sig figs → answer 3 sig figs → 12.5 m/s). Given isotopic data — calculate average atomic mass (worked) (Neon example shown earlier; result 20.180 amu). Using periodic table data to find p, n, e and determine if atom or ion Problem: Given symbol: {}^{37}_{17}\text{Cl}^{-}. Protons = 17 (by atomic number). Mass number = 37 → neutrons = 37 − 17 = 20. Charge −1 → electrons = protons + 1 = 18. This is an anion (ion). Not a neutral atom. Lab hydrate calculations — general example (walkthrough) (We already did CuSO₄·5H₂O example; follow same steps for your lab values.) Density equation & rearrangements (reminder) \rho = \dfrac{m}{V},\qquad m=\rho V,\qquad V=\dfrac{m}{\rho}. Essay-style questions (short, clear answers + reasoning) 1) Classify a substance as element/compound/mixture, pure or mixture, molecules or ions (how to explain) Example substance: Table salt (NaCl) from a bag. Is it pure substance or mixture? If chemically pure NaCl sample → pure substance (compound). If table salt has additives (iodide, anti-caking agents), it’s a mixture. Element/compound? NaCl is a compound (sodium + chlorine chemically bonded). Molecules or ions? Ionic compound made of Na⁺ and Cl⁻ ions (not discrete molecules), so it consists of ions in a crystal lattice. Why: composition fixed for compound; ionic bonding indicates ions rather than molecules. 2) Explain why some elements form cations and others form anions (use noble gas concept) Atoms tend to reach a lower energy, more stable electron configuration. Many atoms achieve stability by adopting the electron configuration of the nearest noble gas: Metals (left side): have few valence electrons; losing them gives a noble gas configuration → form cations (positive). Nonmetals (right side): have more valence electrons but are short of an octet; gaining electrons gives a noble gas configuration → form anions (negative). Example: Na (1 valence electron) loses 1 → Na⁺ (like Ne). Cl (7 valence electrons) gains 1 → Cl⁻ (like Ar). 3) Describe J.J. Thomson and Rutherford experiments (short) Thomson: cathode ray deflection → discovered electrons → atom contains negative particles → plum pudding model. Rutherford: gold foil scattering → most of atom empty; tiny dense positive nucleus deflects alpha particles → nuclear model of atom. Worked examples you specifically asked (Example 1 & 2) Example 1: Using cations: Li⁺ and Ba²⁺; anions: O²⁻ and ClO₄⁻ (perchlorate). Create four neutral ionic compounds and explain subscripts. We want neutral compounds (total positive charge = total negative charge). Possible pairings: Li⁺ + O²⁻ → Li₂O. Reason: O²⁻ has −2, Li⁺ is +1. Need two Li⁺ to balance one O²⁻ → Li₂O. Ba²⁺ + O²⁻ → BaO. Reason: Ba²⁺ (+2) and O²⁻ (−2) already balance 1:1 → BaO. Li⁺ + ClO₄⁻ → LiClO₄. Reason: perchlorate has −1, Li⁺ has +1 → 1:1 ratio. Ba²⁺ + ClO₄⁻ → Ba(ClO₄)₂. Reason: Ba²⁺ (+2) needs two ClO₄⁻ to neutralize → use parentheses for two polyatomic anions. How subscripts determined: cross-balance charges to get net zero; the smallest whole-number ratio is used. Example 2: Write formula or name for each I’ll give the correct formula or name and a brief explanation. K₂S — Name: potassium sulfide. K⁺ and S²⁻ → need two K⁺ to balance one S²⁻. (NH₄)₂SO₄ — Name: ammonium sulfate. NH₄⁺ is +1; sulfate is SO₄²⁻ → need two ammonium to balance sulfate. OBr₂ — Name: oxygen dibromide? Wait — check elements: O + Br is unusual: oxygen normally forms −2, bromine usually −1; a neutral binary molecular compound OBr₂ would be better thought as dibromine monoxide? This is a tricky one. But the formula given is OBr₂; if treated as a binary molecular compound (two elements, both nonmetals), use prefixes: oxygen dibromide (but molecule uncommon). Important note: In practice, more common bromine-oxygen compounds have different formulas (e.g., BrO₂⁻ is chlorite analogs). For classwork, accept oxygen dibromide as name for OBr₂ (prefix naming: mono/di). If the intended was phosphorus triiodide etc., the list might have mixed types; follow molecular naming rules for two nonmetals. Na₂CO₃ · 10H₂O — Name: sodium carbonate decahydrate. Sodium carbonate with ten waters attached. Strontium phosphate — Formula: Sr₃(PO₄)₂. Strontium is Sr²⁺; phosphate is PO₄³⁻. LCM of 2 and 3 is 6 → need three Sr²⁺ (3×+2=+6) and two PO₄³⁻ (2×−3=−6). Gold(III) bromide — Formula: AuBr₃. Gold (III) = Au³⁺, bromide = Br⁻ → need three Br⁻ to neutralize. Beryllium sulfate tetrahydrate — Formula: BeSO₄·4H₂O. Beryllium is Be²⁺, sulfate SO₄²⁻ → BeSO₄, plus 4 waters. Phosphorous triiodide — Formula: PI₃. Binary molecular: phosphorus (P) + iodine (I) → prefixes: phosphorus triiodide. Aluminum hydroxide — Formula: Al(OH)₃. Al³⁺ and OH⁻ → need three OH⁻ to balance one Al³⁺

period 3 n chlorine

ES 3 Manufacturing chlorine

chlorine and carbon

AQA GCSE Chemistry Paper 2

Uses of chlorine

To estimate the concentration of free chlorine in swimming pool water or bleach using: i) A comparator or ii) A colorimeter

Chlorine and Chlorate uses

Part II. Multiple choice, continued (3 points each). Name:__________________ Please circle your answer. There is only one correct answer to each question. 43. The F-N-F bond angle in the NF3 molecule is slightly less than ___________. A) 90˚ B) 109.5 ˚ C) 120˚ D) 180˚ 44. The angles between sp2 orbitals are _____________. A) 90˚ B) 109.5˚ C) 120˚ D) 180˚ 45. The blending of one s atomic orbital and two p atomic orbitals produces __________. A) three sp hybrid orbitals B) three sp2 hybrid orbitals C) three sp3 hybrid orbitals D) two sp3 hybrid orbitals 46. The π bond in ethylene, CH2CH2, results from the overlap of ___________. A) sp3 hybrid orbitals B) s atomic orbitals C) sp2 hybrid orbitals D) p atomic orbitals 47. The amount of gas that occupies 60.82 L at 31˚ C and 367 mm Hg is _________ mol. A) 1.18 B) 0.850 C) 894 D) 11.6 48. The specific heat of lead is 0.13 J/g-K. How much heat (in J) is required to raise the temperature of 15 g of lead from 22 ˚C to 37 ˚C? A) 7.8 B) 1.9 C) 19 D) 29 49. Which one of the following is a weak acid? A) HNO3 B) HI C) HF D) HClO4 50. Which combination will produce a precipitate? A) NH4OH (aq) and HCl (aq) B) AgNO3 (aq) and Ca(C2H3O2)2 (aq) C) NaCl(aq) and HC2H3O2 (aq) D) NaOH (aq) and Fe(NO3)2 (aq) 51. The density of chlorine (Cl2) gas at 25 ˚C and 450 torr is _______ g/L. A) 20 B) 4.9 C) 1.7 D) 0.86

The Reactions of Chlorine

reactions of chlorine and water

Estimate the concentration of free chlorine in swimmingpool water or bleach using a comparator

ES 5 The risks and benefits of using chlorine

Oxoanions and Oxoacids of Chlorine

4. Metals and Non-metals Learning Objectives By the end of the lesson, you will be able to: ☑ distinguish between metals and non-metals ☑ describe the physical and chemical properties of metals and non-metals ☑ list the uses of some metals and non-metals MINERALS AND ORES You have learnt that all materials Here is the exact text from the image:are made up of basic substances called elements, and that elements cannot be split into simpler substances by chemical methods. There are 118 known elements. Sodium, zinc, gold, mercury, iron, lead, barium and tin (metals); and hydrogen, oxygen, carbon, sulphur, chlorine, boron, neon and radon (non-metals) are some examples. Only certain unreactive elements are found free in nature. Others occur in combined states as minerals. A mineral is a solid inorganic substance that is found in nature. A mineral deposit that can be mined and from which an element or compound can be obtained profitably is known as an ore. Elements can be broadly classified into two groups—metals and non-metals. Table 4.1 Some common ores Fig. 4.1 Some common ores a. Bauxite (aluminium) b. Malachite (copper) c. Haematite (iron) d. Galena (lead) e. Apatite (phosphorus) f. Quartz (silicon) -- --- METALS All except 20 of the known elements are metals. Most metals are reactive; they combine with other elements in nature, such as oxygen and sulphur, and occur as oxides, sulphides and carbonates. Only a few unreactive metals like gold, silver and platinum are found as free metals in the Earth's crust. Physical Properties of Metals Metals are solids at room temperature, except mercury, which is a liquid at room temperature (Fig. 4.2(a)). They are generally hard and strong, with a few exceptions such as sodium and potassium, which are soft and can be easily cut with a knife (Fig. 4.2(b)). They have a metallic lustre (shine), especially when freshly cut. They have high melting and boiling points, with a few exceptions like sodium, potassium and mercury. They are good conductors of heat and electricity. Silver and copper are the best conductors of electricity, followed by gold and aluminium. Metals are sonorous. They produce a ringing sound when struck. Most metals have high tensile strength. They can take heavy loads without breaking. They are malleable. Metals, with exceptions like sodium and potassium, can be beaten into thin sheets and foils. They are ductile. Metals, with exception like sodium and potassium, can be drawn into wires. Most metals have high density. However, sodium and potassium have low density and float on water. Fig. 4.2 Special metals a. Mercury b. Sodium --- Chemical Properties of Metals Reaction with oxygen Metals react with oxygen under different conditions to form basic oxides. These basic oxides react with water to form bases. Sodium and potassium react vigorously with oxygen at room temperature. 4Na + O_2 \rightarrow 2Na_2O To prevent this oxidation, sodium and potassium are stored under kerosene. Magnesium reacts with oxygen only when ignited. It burns with a dazzling bright flame and forms a white powder of magnesium oxide. 2Mg + O_2 \rightarrow 2MgO Copper and iron react with oxygen only when heated to a very high temperature. 2Cu + O_2 \rightarrow 2CuO --- --- Reaction with water Metals react with water to form hydroxides or oxides, along with hydrogen. Different metals react at different temperatures. Sodium, potassium, and calcium react with cold water to form hydroxides. 2Na + 2H_2O \rightarrow 2NaOH + H_2 Magnesium Reacts with steam or hot water to form magnesium oxide. Mg + H_2O \rightarrow MgO + H_2 Aluminium Forms an oxide too, but this oxide forms a protective covering over the metal and prevents further reactions. 2Al + 3H_2O \rightarrow Al_2O_3 + 3H_2 Zinc Reacts only with steam. Zn + H_2O \rightarrow ZnO + H_2 Iron Reacts with steam when heated strongly. 2Fe + 3H_2O \rightarrow Fe_3O_4 + 3H_2 Copper, gold, silver, and platinum do not react with water at all. --- Activity 4.1 Teacher Demonstration Aim: To study the reaction of metals with water. [Caution: This activity should be demonstrated by the teacher, and students should stand away from the table.] Materials required: Two 200 mL beakers Pieces of sodium and calcium Forceps Knife Litmus papers Water Method: 1. Fill each beaker with 100 mL of water. 2. Using forceps and a knife, cut a small piece of sodium. 3. Dry it on a tissue paper and drop it into one of the beakers. 4. Repeat the same procedure with calcium. 5. Test the water in both the beakers with red and blue litmus papers. Observations and Conclusions: Sodium reacts vigorously and may explode. A gas is also released. The reaction with calcium is quick, though not as vigorous as that with sodium. In both cases, the red litmus paper turns blue, showing that the solutions are bases. --- Reaction with dilute acids Most metals react with dilute acids to form their salts and liberate hydrogen gas. The reaction with reactive metals like sodium, potassium, and calcium is violent. Magnesium, aluminium, zinc, and iron do not react violently. Mg + 2HCl \rightarrow MgCl_2 + H_2 Copper, silver, gold, and platinum do not react with dilute acids. --- Reaction with bases Only some metals such as aluminium and zinc react with strong bases like sodium hydroxide to liberate hydrogen gas. Zn + 2NaOH \rightarrow Na_2ZnO_2 + H_2 --- Activity 4.2 Aim: To study the reaction of metals with dilute hydrochloric acid. Materials required: Sandpaper Six test tubes Dilute hydrochloric acid Strips of magnesium, zinc, iron, tin, lead, and copper Method: 1. Clean the metal strips with sandpaper. 2. Add dilute hydrochloric acid to the six test tubes. 3. Insert a strip of metal into each test tube. Observe if any bubbles are formed in the test tubes. If no bubbles are seen, warm them gently in a beaker of hot water. 4. Observe the speed at which gas is generated. This gives an idea of the speed of the reaction. 5. Classify the metals in order of their reactivity with dilute hydrochloric acid. [Caution: Acids are corrosive and should be handled carefully.] --- Activity 4.3 Aim: To study the reaction of metals with bases. Materials required: Small piece of zinc Beaker Sodium hydroxide Method: 1. Prepare warm sodium hydroxide or caustic soda solution. 2. Drop the piece of zinc into it. Observations and Conclusions: You will notice that zinc reacts with sodium hydroxide to liberate hydrogen gas. Observations on Metals with Dilute Acids Metals like sodium, potassium, and calcium react violently with dilute acids to liberate hydrogen gas. Magnesium, aluminium, zinc, and iron also displace hydrogen from dilute acids, but the reaction is not violent. Metals such as copper, silver, gold, and platinum do not displace hydrogen from dilute acids. --- Activity Series of Metals The activity series of metals is the arrangement of metals in decreasing order of reactivity. The series in the book shows reactivity decreasing from top to bottom. Potassium is the most reactive metal while gold is the least reactive. --- Displacement of a Metal by Other Metals A more reactive metal displaces a less reactive metal from its compounds in an aqueous solution. Some examples: Mg + CuSO_4 \rightarrow MgSO_4 + Cu Zn + FeSO_4 \rightarrow ZnSO_4 + Fe Iron can displace copper from copper sulphate solution (as shown in Activity 4.4). The solution turns green, and reddish-brown copper deposits on the iron nail. Copper cannot displace iron from iron sulphate solution, showing that copper is less reactive than iron. Cu + FeSO_4 \rightarrow \text{No reaction} Question: What do you think will happen if you place a silver spoon in copper sulphate solution? --- Activity 4.4 - Displacement Reaction Aim: To study a displacement reaction. Materials Required: Test tube Iron nail Copper sulphate solution Method: 1. Fill the test tube with copper sulphate solution (blue in colour). 2. Place the clean iron nail in the solution. Observations and Conclusions: After about an hour, the solution changes to green, and a reddish-brown deposit is formed on the iron nail. --- Corrosion of Metals Corrosion is the destruction or damage of a material due to chemical reaction. Rusting of iron happens when iron is exposed to moist air, forming a reddish-brown layer of rust. Rust is iron oxide, which eventually flakes off, damaging the object. Definition written on the page: "Slow eating of a metal’s surface due to oxidation is called corrosion of metals." --Observations on Metals with Dilute Acids Metals like sodium, potassium, and calcium react violently with dilute acids to liberate hydrogen gas. Magnesium, aluminium, zinc, and iron also displace hydrogen from dilute acids, but the reaction is not violent. Metals such as copper, silver, gold, and platinum do not displace hydrogen from dilute acids. --- Activity Series of Metals The activity series of metals is the arrangement of metals in decreasing order of reactivity. The series in the book shows reactivity decreasing from top to bottom. Potassium is the most reactive metal while gold is the least reactive. --- Displacement of a Metal by Other Metals A more reactive metal displaces a less reactive metal from its compounds in an aqueous solution. Some examples: Mg + CuSO_4 \rightarrow MgSO_4 + Cu Zn + FeSO_4 \rightarrow ZnSO_4 + Fe Iron can displace copper from copper sulphate solution (as shown in Activity 4.4). The solution turns green, and reddish-brown copper deposits on the iron nail. Copper cannot displace iron from iron sulphate solution, showing that copper is less reactive than iron. Cu + FeSO_4 \rightarrow \text{No reaction} Question: What do you think will happen if you place a silver spoon in copper sulphate solution? --- Activity 4.4 - Displacement Reaction Aim: To study a displacement reaction. Materials Required: Test tube Iron nail Copper sulphate solution Method: 1. Fill the test tube with copper sulphate solution (blue in colour). 2. Place the clean iron nail in the solution. Observations and Conclusions: After about an hour, the solution changes to green, and a reddish-brown deposit is formed on the iron nail. --- Corrosion of Metals Corrosion is the destruction or damage of a material due to chemical reaction. Rusting of iron happens when iron is exposed to moist air, forming a reddish-brown layer of rust. Rust is iron oxide, which eventually flakes off, damaging the object. Definition written on the page: "Slow eating of a metal’s surface due to oxidation is called corrosion of metals." Uses of Metals (Continued) Aluminium Used in high-voltage electric lines. Alloys like duralumin and magnalium are used in aircraft and automobile bodies. Used for making aluminium foil and cooking utensils. Copper Good conductor of electricity → Used in electrical wires, cables, motors, and transformers. Good conductor of heat → Used in the bottoms of stainless steel vessels. Zinc Used to make corrosion-resistant galvanised iron (GI) pipes and sheets. Used as an electrode in dry cells. Other Metals Gold and silver → Used in jewellery. Lead → Used in electrodes of lead storage batteries (used in automobiles and inverters). Chromium → Used for electroplating iron to give a shiny, corrosion-resistant finish. --- Looking Back (True/False Statements) 1. Gold, silver, and platinum are found in the Earth’s crust as free metals. → True 2. Most metals are solids that are soft. → False 3. Metals such as zinc and magnesium react with dilute acids to liberate oxygen. → False 4. A less reactive metal displaces a more reactive metal from its aqueous solution. → False 5. The chemical name of rust is zinc oxide. → False (Rust is Fe₂O₃.xH₂O) 6. Coating zinc objects with iron is called galvanising. → False (Galvanising is coating iron with zinc) Non-Metals Physical Properties of Non-Metals Exist as gases or solids at room temperature (except bromine, which is liquid). Not as hard as metals (except diamond, which is very hard). Low tensile strength and low density. Low melting and boiling points (except graphite). Not sonorous (do not produce a ringing sound). Not malleable or ductile (cannot be beaten into sheets or drawn into wires). Do not have lustre (except iodine and graphite). Bad conductors of heat and electricity (except graphite, and silicon under specific conditions). --Chemical Properties of Non-Metals Reaction with Water Most non-metals do not react with water. Highly reactive non-metals (e.g., phosphorus) catch fire in air, so they are stored in water. Fluorine, chlorine, and bromine react with water to form acids. Reaction with Oxygen Non-metals react with oxygen to form acidic or neutral oxides. Carbon and sulfur react with oxygen to form acidic oxides, which dissolve in water to form acids. Some oxides (e.g., CO, N₂O) are neutral and do not form acids. Examples: Carbon + Oxygen → Carbon Dioxide (CO₂) CO₂ + Water → Carbonic Acid (H₂CO₃) Sulfur + Oxygen → Sulfur Dioxide (SO₂) SO₂ + Water → Sulfurous Acid (H₂SO₃) Reaction with Acids Unlike metals, non-metals do not replace hydrogen in acids. Silicon reacts with hydrofluoric acid (HF). --Uses of Non-Metals Hydrogen Used in the manufacture of ammonia and industrial chemicals. Used in vanaspati (a cooking oil). Oxygen Used in breathing support systems in hospitals. Used with other gases in equipment to weld metals. Sulphur Used in the manufacture of sulphuric acid, sulphur dioxide gas, and other industrial chemicals. Used to make pesticides for agriculture. Used in vulcanising rubber (making it harder) and in gunpowder. Nitrogen Used in the manufacture of ammonia and nitrogenous fertilisers like ammonium nitrate and ammonium sulphate. Used as an inert gas in processed food packaging to prevent rancidity. Silicon Used in making semiconductors for microchips. Silicates (oxides of silicon) are used in making glass. Other Non-Metals Phosphorus: Used in making fertilisers (superphosphates). Chlorine: Used for disinfecting drinking water. Argon: Used in welding stainless steel and filling electric bulbs. Helium: Used in balloons for meteorological observations. Neon: Used in fluorescent lights for advertisement displays

chlorine and carbon polyatomic

Chlorine Common Polyatomic Ions

Pathogens, Chlorine, and UV Disinfection Quiz

B) Group 7 (halogens) – chlorine, bromine and iodine