SnPBlock-Lecture T2-2024 PDF

The s-Block Elements  Members of the s-Block Elements I II I- Alkali metals Li Be II-Alkaline Earth metals Na Mg constitute K Ca the s-block Rb Sr elements Cs Ba Fr Ra ...

The s-Block Elements  Members of the s-Block Elements I II I- Alkali metals Li Be II-Alkaline Earth metals Na Mg constitute K Ca the s-block Rb Sr elements Cs Ba Fr Ra  Group I and II elements Their outermost shell electrons are in the s sub shell Elements in the same group also undergo similar chemical reactions Group I elements Group II elements Lithium Beryllium Sodium Magnesium Potassium calcium Rubidium Strontium Barium Silvery ,white metals ,exposing a shiny surface that tarnishes Francium – radioactive rapidly in air due to oxidation.  The s-Block Elements Similarities 1. highly reactive metals 2. strong reducing agents 3. form ionic compounds with fixed oxidation states +1 for Group I elements +2 for Group II elements  Properties of the s-Block Elements Trends that can be analysed down the group include atomic radius first ionization energy melting point and reactivity. Atomic Radius Group I Atomic radius Group II Atomic element (nm) element radius (nm) Li 0.152 Be 0.112 Na 0.186 Mg 0.160 K 0.231 Ca 0.197 Rb 0.244 Sr 0.215 Cs 0.262 Ba 0.217 Fr 0.270 Ra 0.220 Down the groups atomic radius increases the outermost electrons are further away from the nuclei Atomic radius of group II < Group I Size decreases from left to right across the periods  Ionization Energy The metallic radii and the value of the 1st IE for the group 1 elements. Ionisation energy (IE) decreases down the groups. The energy needed for the ionization is used to overcome the electrostatic attraction between the electron being removed and the protons in the nucleus. Down the group Nuclear charge (the number of protons in the nucleus) increases However, the number of filled inner shells increases (shielding/screening effect increases ie: the repulsion between filled inner shells and the electron being removed increases) The outer electrons added to a new quantum shell, further from the nucleus and more shielded Attraction between the nucleus and the valence electron weakens causing a decrease in IE.  Melting point depends on the Melting point strength of its metallic bonds Bonding: Strength of metallic bond : Group II > Group I decreases down the group because the atomic radius increases, resulting in a weaker attraction between the nucleus and delocalized electrons Low melting points m.p./b.p. : Group II > Group I Density Density : Usually low density Generally increases down the group Group II > Group I  Metallic character refers to chemical properties that are associated with the elements classified as metals which depends on the ability of an element to readily lose its outer valence electrons to form positive ions. Ionization energy decreases → easy to lose electrons → more reactive. Atomic radius increases → outer electrons further away from the nucleus → Atomic size decreases electron less Ionisation energy increases attracted by the Atoms more readily accept nucleus. electrons to fill valence shell  Characteristics of s-block elements soft, lustrous, reactive metals High tendency to lose e- to form positive ions Metallic character Group I more reactive than increases down both group II groups. Hardness decreases down the group Group I < Group II Na/K…can be easily cut with a knife  Reactions with WATER Group I: They all react vigorously with water 2Li (s) + 2 H2O(l) 2LiOH(aq) + H2(g) 2Na (s) + 2 H2O(l) 2NaOH(aq) + H2(g) 2K (s) + 2 H2O(l) 2KOH(aq) + H2(g) 2Rb (s) + 2 H2O(l) 2RbOH(aq) + H2(g) The hydroxides of the gp I are among the strongest bases https://www.youtube.com/watch?v=jI__JY7pqOM  Lithium Sodium Potassium https://www.youtube.com/watch?v=Vxqe_ZOwsHs https://www.youtube.com/watch?v=dmcfsEEogxs https://www.youtube.com/watch?v=oqMN3y8k9So https://www.youtube.com/watch?v=0YNsIaSbFdg Reactions with water or steam Group II – less reactive than Group I Be has no reaction with either water or steam even at red heat Reaction of Mg with water is a very slow (does not proceed completely) a reaction can be seem with WARM water Mg(s) + H2O(g) MgO(s) + H2(g) STEAM Mg(s) + 2H2O(l) → Mg(OH)2(s) + H2(g) This is a much slower reaction than the reaction with steam Ca with water Calcium, Strontium and barium reacts vigorously with cold water with increasing vigour down the group to form hydroxides. https://www.youtube.com/watch?v=-gu1vlJ28co https://www.youtube.com/watch?v=uwIysB3y3fo  https://www.youtube.com/watch?v=i-rFsFwdkTU Phenolphthalein indicator turn from colourless to pink with alkalis https://www.youtube.com/watch?v=B2ZPrg9IVEo Reactivity increases down the group can Reaction of Group II be seem by the increase in effervescence metals with water Barium reacts very rapidly with water Magnesium reacts with cold water extremely Strontium reacts slow fairly quickly with cold water  Increases down Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g) the group Sr(s) + 2H2O(l) Sr(OH)2(aq) + H2(g) Ba(s) + 2H2O(l) Ba(OH)2(aq) + H2(g) The hydroxides produced make the water alkaline observation: fizzing, (more vigorous down group) the solution heating up (more down group) and with calcium a white precipitate appearing (less precipitate forms down group) Down the group 2, the resulting solutions become more alkaline because the concentration of hydroxide ions (OH-) in the solution formed increases Complete equations for the reactions given on page 2. Reactions with OXYGEN Metals in group I and II reacts vigorously with oxygen Most s-block elements -show a silvery white lustre Sodium shows a silvery white when they are freshly cut lustre when freshly cut -they tarnish (form a dull dark layer on the surface) rapidly upon exposure to the oxygen in the air. -they react with oxygen in the air to form an oxide layer Na can forms peroxides and K ,Rb ,Cs form super oxides. Li(s) + O2(g) → K(s) + O2(g) → Rb(s) + O2(g) → Lithium is the LEAST reactive Group I metal and the reactivity of the metals increases down Group 1 Reaction with oxygen Mg will react slowly with oxygen without a flame. Mg ribbon will often have a thin layer of magnesium oxide on it formed by reaction with oxygen Mg burns with a bright white flame.  S-block elements reacts readily with oxygen to form solid metal oxides. Ca(s) + O2 (g) → Sr(s) + O2 (g) → The elements need to be heated for the reaction to start. Without heating, a slow reaction between the element and oxygen when exposed to air. Surface coating of oxide prevents the element from further reaction. Barium is the MOST reactive and often stored under oil  RUBIDIUM CAESIUM Most of them have to be stored under liquid paraffin to prevent contact with the atmosphere. Reaction with chlorine Reaction of metals with dilute acids All group I and II metals react with acids to give hydrogen gas and the corresponding salt. Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g) The reactivity of group II metals with acid increases down the group. Be reacts slowly with acids. Mg, Ca, Sr and Ba reacts vigorously with acids.  Explaining trend in reactivity of group I and II The reactivity of s-block elements INCREASES down the groups. Ionization energy decreases down the group and increases across a period Less energy required to remove outer electrons - nuclear charge increases - Greater shielding effect by extra inner shells of electrons - larger distance of outermost electrons from the nucleus outweigh the attraction of higher nuclear charge (The increase in the atomic radius results in a decrease in the attractive force between the outer electrons and the nucleus) These cause a decrease in energy needed to remove the electron. Red litmus turns blue in alkalis less basic than Group I  Solubility of Group II hydroxides Compound Solubility / mol per 100g water Solubility of hydroxides Be(OH)2 insoluble increases down Mg(OH)2 very slightly soluble the group. Ca(OH)2 slightly soluble As the solubility of the group 2 hydroxides increases, so Sr(OH)2 quite soluble does the pH of the solutions Ba(OH)2 Very soluble formed. This is because the more of Mg(OH)2 - Milk of Magnesia the hydroxide that dissolves, the greater the concentration Ba(OH)2 - Not safe to drink of hydroxide ions (OH-) in the solution formed. Mg2+(aq) + 2OH-(aq) → Mg(OH)2(s) Ba2+(aq) + 2OH-(aq) → Ba(OH)2(aq)  Hydroxides Group I - Hydroxides All group I LiOH Solutions become more basic oxides/hydroxides NaOH as the solubility of the are soluble and KOH hydroxides increases basic (alkaline) RbOH CsOH Group II - Hydroxides Group II hydroxides are less Solubility Be(OH)2 soluble and therefore the increase, from Mg(OH)2 oxides form weaker alkaline amphoteric to Ca(OH)2 solutions than those of basic, base group I. Sr(OH)2 strength increase Ba(OH) Beryllium oxide/hydroxide 2 are amphoteric.  The compounds are all ionic except BeO which has covalent character. All these oxides are basic in nature (except beryllium oxide which is amphoteric) basic – reacts with acids to give salt and water Na2O(s) 2 HCl → NaCl(aq) + H2O(l) MgO (s) + 2 HCl (aq) → MgCl2 (aq) + H2O (l) SrO (s) + 2 HCl (aq) → SrCl2 (aq) + H2O (l) amphoteric – reacts with both acids and bases Be(OH)2(s) + 2HCl(aq) → BeCl2(aq) + 2H2O(l) Be(OH)2(s) + 2NaOH (aq) → Na2Be(OH)4 (aq) sodium tetrahydroxoberyllate(II)  Solubility of Group II sulfates Compound Solubility / mol per 100g water BeSO4 very soluble Solubility of sulfates decreases down the MgSO4 quite soluble group. CaSO4 slightly soluble SrSO4 very slightly Note that this decrease in soluble/insoluble solubility down the group BaSO4 insoluble is the opposite of the trend for the solubility of the BaSO4 - X- rays are blocked group 2 hydroxides. Mg2+(aq) + SO42-(aq) → MgSO4(aq) Ba2+(aq) + SO42-(aq) → BaSO4(s) Barium sulphate is used as a radiocontrast agent to help take X-ray images of the digestive system. It is sometimes known as a ‘barium meal’. Barium sulphate is insoluble, so is not absorbed by the body when swallowed. However, barium is a very good absorber of X-rays and it helps to define structures of the digestive system to aid in diagnosis.  Solubility of the Sulfates and Hydroxides of s-Block Elements M+ Group I >> Group II All Group 1 hydroxides and sulfates are soluble. Not all group 2 hydroxides and sulfates are soluble. Group I compounds are more soluble than Group II because the metal ions have smaller charges and larger sizes. Group 2 cations are doubly charged and have smaller sizes – less soluble than Group I.  Thermal stability Why do some compounds decompose when heated while others don't?  Thermal stability refers to decomposition of the compound on heating. A compound is often said to be thermally stable if it does not decomposed at the temperature of the normal Bunsen flame (approximately 1000oC). Increased thermal stability means a higher temperature is needed to decompose the compound. Thermal stability of nitrates and carbonates Group I and II, carbonates and nitrates when heated they decompose. NO3- The stabilities of nitrate and carbonates anions are influenced by the CHARGE and SIZE of the cation present carbonate ion (CO32-)  The thermal stability of an ionic compound depends on Polarizing power of cation 1+ 2+ Li Be Cations with small size and high charge has a higher polarising power They can polarise the anion more Na Mg effectively. This assists decomposition and makes K Ca the compound less stable to heat. Charge density DECREASES The large anion is polarized by the small cation Thermal stability INCREASES on DOWN the group. Metal ions become larger down group but have the same charge. This means their charge density is reduced. The compound less stable more likely to decompose on heating When a compound with large anions undergoes thermal decomposition, a compound with small anions will be formed since small anions are less easily polarized Observations obtained on heating samples of nitrates No brown fumes indicate a lesser decomposition metal nitrate → metal nitrite + oxygen gas If brown fumes observed, it indicates a greater decomposition metal nitrate →metal oxide + NO2 + O2 colourless brown gas  Thermal stability DECREASES Group I nitrates are white solids. Group I nitrates, except LiNO3, decomposes to form metal nitrites and oxygen gas. Metal Nitrite + O2 Thermal stability INCREASES LiNO3 decomposes in a similar manner to group II nitrates  Thermal decomposition of group II metal nitrates forms the metal oxide, nitrogen dioxide and oxygen. 2M(NO3)2(s) → 2MO(s) + O2(g) + 4NO2(g) brown gas Decomposition temperature magnesium nitrate: Mg(NO3)2 nitrates become increases (more more stable to heat energy calcium nitrate: Ca(NO3)2 thermal needed) strontium nitrate: Sr(NO3)2 decomposition increasing stability barium nitrate: Ba(NO3)2 down the group. + - + - Polarization of electron cloud o + - + - + - Polarization of electron cloud of an anion, by cation.  - + - ectron cloud of an anion, by cation.  O M2+ -O N O  O M2+ -O N O N-O bond is polarised  O M2+ -O N O  O M2+ -O N O M2+ O2- + NO2  Decomposition of Group II nitrates occurs Bond is weakened due to polarisation of electron cloud by cation O Heat -O N + NO2 M2+ M2+ O2- O  Experiments to study Thermal Decomposition of carbonates Place the same amount of carbonate in a glass test tube. Fix a delivery tube to the test tube and clamp in a stand, as in the figure. Light the Bunsen and measure the time taken for the gas evolved to change the color of limewater Group I carbonates are Thermal stability DECREASES stable to heat. They do not decompose at Bunsen temperatures except lithium carbonate, All Groups II metals carbonates which decomposes in a undergo decomposition on heating similar manner to group II carbonates. No reaction Thermal stability INCREASES stable to heat down group 1 and 2 Lithium carbonate is relatively unstable as lithium ion has the smallest size and the polarizing power is highest The electron cloud of any large anion is distorted to a great extent and decompose more readily on heating Group II metal ions O are smaller in size Mg2+ - :O C Mg2+ O2- + CO2 have higher charge O:- therefore stronger polarizing power The electron cloud of the carbonate ion is much distorted, therefore more readily undergo thermal decomposition MgCO3 MgO The oxide ions are smaller in size than carbonate ions so they are less polarizable Lithium carbonate and Group II carbonates on heating decompose to give oxides upon heating BaCO3 BaO Lithium and group 2 carbonates decompose when heated to form the metal oxide and carbon dioxide. The group 2 carbonates become more stable to thermal decomposition going down the group: Decomposition temperature increases (more heat energy needed) increasing MgCO3 stability CaCO3 SrCO3 BaCO3 decomposes but at a Harder to decompose BaCO3 very high temperature  The polarising 2 power of the metal cation decreases down the group as 2 the charge-density decreases The carbonate anion has a relatively large ionic radius so it is easily polarized by a small highly charged cation polarization increasing MgCO3 DOWN Charge density DECREASES THE Polarizing power of cation DECREASES stability CaCO3 GROUP Harder to SrCO3 decompose BaCO3 BaCO3 do not decompose at Bunsen temperatures Effect of sizes of the cations on thermal stability of the carbonates of both Groups I and II metals Thermal stability of group I and II carbonates across the period decreases because  As you go down a group, the atom + - gets bigger, hence the positive ion Decreasing gets bigger and so polarizing they have less - polarising effect on power + the carbonate / nitrate ions near them. Therefore as you + - go down the group, the group 1 and Increasing group 2 nitrates stability and carbonates become more thermally stable. In summary, down the group Charge Polarising Size of cation density of power of increases cation cation decreases decreases Decomposition Thermal stability temperature increases increases In summary, down the group Size of cation Thermal stability increases increases Decomposition temperature increases All group 2 metal salts decomposes on heating. Thermal stability increases down the group. All group 1 metals are more stable to heat except Lithium salts  Some Applications of thermal stability Cement and Lime Production Calcium carbonate (CaCO₃) Raw material in the production of cement and quicklime (CaO): CaCO₃ → CaO + CO₂ MgCO₃ and CaCO₃ are used in the production of glass and ceramics: – MgO, CaO (obtained from decomposition) required for forming the structure of glass and ceramic products. – MgO, which is used as a key ingredient in antacids.  P-Block GROUP VII - The Halogens  ‘hal’ → salt ‘gen’ → produce Halogens react with metals to form metal salts The term ‘halogen’ is derived from Greek meaning ‘salt making’. Halogens are very reactive and readily form salts. Electron configuration All have the electronic configuration…..ns2 np5. Electrons go into shells further from the nucleus Properties F Cl Br I Molecular formula F2 Cl2 Br2 I2 Bonding Covalent Covalent Covalent Covalent They are all non-metals and exist has separate diatomic molecules. They are non-metals and have electronic configurations just one electron short of the nearest noble gas  Properties F2 Cl2 Br2 I2 Colour Yellow Green Red/brown Grey State (at RTP) GAS GAS LIQUID SOLID Vapour colour Yellow Green Red/brown Purple The halogens exist as diatomic molecules that interact through dispersion forces, which increase in strength as the atoms become larger. 2 non-polar molecules Polarity in one molecule Induced dipole  Physical properties Fluorine (F2): very pale yellow gas. It is highly reactive Chlorine : (Cl2) greenish, reactive gas, poisonous in high concentrations Bromine (Br2) : red liquid, that gives off dense brown/orange poisonous fumes Iodine (I2) : shiny grey solid sublimes to purple gas.  Trends in melting and boiling point Down the group, This leads to greater melting and boiling London forces between temperatures molecules, increasing the increases. energy needed to separate the molecules and therefore higher melting and boiling points. Strength of London forces increases as London the number of forces electrons increases fluorine iodine boiling point = -118 °C boiling point = 184 °C  When bromine is left at room temperature, it gives off brown vapour, as its boiling point (59oC) is not much higher than room temperature. Br2(l) Br2(g) When iodine is warmed, most of it changes directly into a vapour without melting. This change is called sublimation. I2(s) I2(g) Halogens become LESS volatile because as the size of halogen molecules increases as the number of electrons increase. This increases the strength of London forces. Therefore more energy is required to overcome the forces of attraction between molecules.  Atomic and ionic radius down the group Atomic radius / nm Ionic radius / nm F 0.064 0.136 F¯ Cl 0.099 0.181 Cl¯ Br 0.111 0.195 Br¯ I 0.128 0.216 I¯ ATOMIC RADIUS/ IONIC RADIUS INCREASES down the group because the number of shells increases. Ionic radius is larger than the atomic radius, because, the effective nuclear charge decreases with the addition of an extra electron to the outer shell. Repulsion exists between the electrons and the added electrons.  Electron affinity 1st electron affinity F -348 Cl -364 Br -342 I -314 Electron affinity generally decreases down the group (from chlorine to iodine). This is because as the size of the atom increases, the attraction between the incoming electron and the nucleus decreases.  Electronegativity Electronegativity F 4.0 Electronegativity is the ability of Cl 3.0 an atom to attract the pair of Br 2.8 I 2.5 electrons in a covalent bond. Electronegativity decreases down the group. As the number of shells increases, the shielding of the nucleus by inner shell electrons increases and the size of the atom increases so the ability of the nucleus to attract electron pair is reduced. Polar Non Polar  Trends in Reactivity F The reactivity of group 7 elements DECREASES down the group. F2 is the most reactive and I2 the least reactive. Cl F(g) + e- → F-(g) Br Most the reactions of halogens involve them gain electrons to form negative ions (acting as oxidizing agents) I halogens gain one electron to form a negative ion of charge –1) Generally, electron affinity decreases hence it is less spontaneous to form anions (negatively charged ions) OR by gaining slightly negative (δ-) part of the polar molecule. Reacts with other non-metals to form covalent compounds (sharing electron pair with non-metal atom, thus forming a covalent bond). Electronegativity DECREASES down the group Decrease in reactivity is also related with the decrease in electronegativity down the group Fluorine is the most electronegative element of ALL electron configuration: Is2 2s2 2p5 Smaller size (distance between the nucleus and the bonding pair of electrons is very small). Less shielding due to only two electron in the inner shell Power to attract the electron pair is very high The exceptional reactivity of elemental F2 is also related to the weakness of the F-F bond. (Fluorine is an exception due to its extremely small size. The F-F bond length is so short that the lone pairs of electrons on the fluorine atoms repel each other and weakens the F-F bond)  CHEMICAL PROPERTIES F Cl Br I Ion F¯ Cl¯ Br¯ I¯ Configuration 2,8 2,8,8 2,8,18,8 2,8,18,18,8 Reactivity Increasingly reactive Reacts with Metals in Group I and II to produce ionic solids (usually white) reactive NON-METALS Most vigorous between Least vigorous elements at the between bottom of gp I beryllium and and II and top iodine or of gp 7 astatine  Displacement reactions of Halogens Due to high electronegativity, most reactions of the halogens involve them acting as oxidizing agents and gaining electrons to form negative ions or becoming the slightly negative part of a polar molecule. Strong oxidizing agent. They can remove electron from other substances. The oxidizing power of the halogen DECREASES as we move down the group from F to I (they need one electron to complete their octet. They accept electrons and get reduced) F2(g) + 2e- → 2F-(aq) F2 > Cl2 > Br2 > I2 strongest weakest Elements become less reactive (electronegativity decrease) In displacement reactions between halogens and halides, the halogen acts as an oxidizing agent. Decreasing oxidizing ability halide ion: a halogen atom with a negative charge fluorine A displacement reaction is where one species takes the place of another in a compound. chlorine A more reactive halogen will displace a less reactive halogen from its bromine compounds. Halogens get less reactive as it goes down the group. iodine Chlorine displaces bromine and iodine Cl2 + 2NaBr → Br2 + 2NaCl Cl2 + 2NaI → I2 + 2NaCl Bromine displaces iodine but not chlorine Br2 + 2NaCl → NO REACTION Br2 + 2NaI → I2 + 2NaCl Iodine does not displace either chlorine or bromine. Add about 2 cm3 of halogen (eg: chlorine water) Add sodium bromide (salt) to chlorine water.  Difficult to differentiate between iodine solution and bromine solution The two solutions can easily be distinguished by adding some hydrocarbon solvent (cyclohexane) and shaking the mixture. Add about 2 cm3 of halogen (eg: chlorine water) Add potassium bromide (salt) to chlorine water. Add about 2 cm3 of cyclohexane Shake and observe. the trend can be explained by considering the nucleus’s attraction for the incoming electron which is affected by the... INCREASED SHIELDING EFFECT INCREASING ATOMIC RADIUS The incoming electron gets further from the nucleus and less attracted by the protons in the nucleus More number of complete inner shells (energy levels) of electrons, so incoming electron experiences MORE repulsion Fluorine Chlorine Bromine Iodine Very Strong Fairly Mild strong oxidising strong oxidising oxidising agent oxidising agent agent agent Oxidising power decreases down the group Electrons are gained less readily  Reaction with Hydrogen Direct combination Hydrogen halides can be made by direct combination H2(g) + F2(g) → 2HF(g) explodes with hydrogen under all conditions H2(g) + Cl2(g) → 2HCl(g) explodes in direct sunlight, proceeds slowly in dark H2(g) + Br2(g) → 2HBr(g) reacts with hydrogen slowly on heating H2(g) + I2(g) 2HI(g) No reaction unless strongly heated, proceeds slowly and only partially  Hydrogen halides as ACIDS colourless gases, polar diatomic molecules HX(g) + H2O(l) → H3O+(aq) + X-(aq) The hydrogen halides (HX molecules) readily react with water to form a hydrohalic acid solutions which are colourless. Strong acids are acids that are completely or nearly 100% ionized in their solutions. The extent to which dissociation occurs is largely depends on the HX bond strength. weakest acid strongest acid HF, with its relatively short, strong bond forms a weak acid. HF(g) + H2O(l) H3O+(aq) + F-(aq) The others completely dissociate HBr(g) + H2O(l) → H3O+(aq) + Br-(aq) Acid strength of hydrogen halides (HX) HI > HBr > HCl >> HF HF is a weak acid because H-F bond is very strong Fluorine has the highest electronegativity  Testing for halide ions in solution Halides can be identified by their reaction with acidified silver nitrate solution to form silver halide precipitates. Dissolve a small amount of halide compound (salt) in water acidify with dilute nitric acid – this prevents the precipitation of other salts (remove CN-, CO32-, SO32- ) This test cannot be used Add a few drops of silver detect fluoride, F- ions nitrate solution, AgNO3. in aqueous solutions, as AgF is soluble AgCl AgBr AgI  The silver ions, Ag+ combines with the halide ions, X- to form a silver halide precipitate. The silver halide precipitates are coloured depending upon the halide present (possible to identify the halide) Potassium Chloride + Silver nitrate KCl(aq) + AgNO3(aq) → KNO3(aq) + AgCl(s) white precipitate Ionic: Cl–(aq) + Ag+ (aq) → AgCl(s) Treat any precipitate with dilute ammonia solution if a precipitate still exists, add concentrated ammonia solution. This can then be used to confirm the result. CHLORIDE white ppt of AgCl soluble in dilute ammonia BROMIDE cream ppt of AgBr insoluble in dilute ammonia but soluble in conc NH3 IODIDE yellow ppt of AgI insoluble in dilute and conc. ammonia solution KCl(aq) + AgNO3(aq) → KNO3(aq) + AgCl(s) Silver chloride has a low solubility in water, so it forms a white precipitate. The positive result in the test for chloride ions.