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Group 1 forms 1+ ions
Group 2 forms 2+ ions

","mcqOptions":[{"text":"

Group 1 forms 1+ ions
Group 2 forms 2+ ions

","correct":true,"explanation":"Group 1 atoms lose one outer electron to form 1+ ions, while Group 2 atoms lose two outer electrons to form 2+ ions."},{"text":"

Group 1 forms 2+ ions
Group 2 forms 1+ ions

","correct":false,"explanation":"This reverses the charges associated with the two groups. Group 1 metals lose one electron to form 1+ ions, whereas Group 2 metals lose two to form 2+ ions."},{"text":"

Group 1 forms 1- ions
Group 2 forms 2- ions

","correct":false,"explanation":"Negative ions form by gaining electrons, but these metals lose their outer electrons. Group 1 therefore forms 1+ ions and Group 2 forms 2+ ions."},{"text":"

Group 1 forms 2+ ions
Group 2 forms 3+ ions

","correct":false,"explanation":"This adds one to each charge by miscounting the outer electrons lost. Group 1 forms 1+ ions, while Group 2 forms 2+ ions."}]},{"id":"04f59550-0241-49c7-b793-4c3e8a38e757","cardType":"cloze","points":null,"front":"First ionisation energy is the energy required to remove one mole of electrons from one mole of [ ] to form one mole of gaseous 1+ ions.","back":"First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.","mcqOptions":[{"text":"gaseous atoms","correct":true,"explanation":"Ionisation energy is defined for isolated gaseous atoms, ensuring that the measured energy concerns electron removal without additional energy changes from separating particles."},{"text":"solid atoms","correct":false,"explanation":"Using solid atoms would include energy associated with overcoming metallic bonding. First ionisation energy specifically removes electrons from one mole of gaseous atoms."},{"text":"gaseous molecules","correct":false,"explanation":"Molecules would involve molecular bonding and could produce molecular ions. First ionisation energy is defined using one mole of isolated gaseous atoms."},{"text":"aqueous ions","correct":false,"explanation":"Aqueous ions are already charged and interact with water molecules, so they do not fit the definition. The starting particles must be gaseous atoms."}]},{"id":"5257236b-a80c-4e21-8ded-570d3966a353","cardType":"cloze","points":null,"front":"First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of [ ].","back":"First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.","mcqOptions":[{"text":"gaseous 1+ ions","correct":true,"explanation":"Removing one electron from each gaseous atom produces one mole of gaseous ions carrying a single positive charge, so gaseous 1+ ions are formed."},{"text":"gaseous 2+ ions","correct":false,"explanation":"Forming 2+ ions requires removing two electrons per atom and relates to two successive ionisations. First ionisation forms gaseous 1+ ions."},{"text":"gaseous 1- ions","correct":false,"explanation":"A 1- ion forms when an atom gains an electron, not when it loses one. First ionisation therefore produces gaseous 1+ ions."},{"text":"aqueous 1+ ions","correct":false,"explanation":"The ions do carry a 1+ charge, but they remain in the gaseous state in the definition. Dissolving them would introduce hydration energy changes."}]},{"id":"dc550126-889d-463f-ae13-681d25a31b63","cardType":"cloze","points":null,"front":"The first ionisation energy of Group 1 and 2 metals [ ] down the group.","back":"The first ionisation energy of Group 1 and 2 metals decreases down the group.","mcqOptions":[{"text":"decreases","correct":true,"explanation":"Down each group, increased atomic radius and electron shielding weaken the attraction between the nucleus and the outer electron, so first ionisation energy decreases."},{"text":"increases","correct":false,"explanation":"Although nuclear charge increases down the group, greater distance and shielding outweigh it. The outer electron is removed more easily, so first ionisation energy decreases."},{"text":"remains constant","correct":false,"explanation":"Each successive element has an additional electron shell, changing both radius and shielding. These effects weaken nuclear attraction, so first ionisation energy decreases rather than remaining constant."},{"text":"doubles","correct":false,"explanation":"Ionisation energy does not follow a simple multiplication pattern down a group. Increased radius and shielding make electron removal easier, so the value decreases."}]},{"id":"01a6fa45-b202-411b-af73-e153a63a9201","cardType":"basic","points":3,"front":"Explain why the reactivity of Group 1 and Group 2 metals increases down the group.","back":"

Atomic radius increases
Electron shielding increases
Attraction between nucleus and outer electrons decreases

","mcqOptions":[{"text":"

Atomic radius increases
Electron shielding increases
Attraction between nucleus and outer electrons decreases

","correct":true,"explanation":"Additional electron shells increase radius and shielding down the group. These effects weaken attraction to the outer electrons, making electron loss and metal reactions easier."},{"text":"

Atomic radius decreases
Electron shielding decreases
Attraction between nucleus and outer electrons increases

","correct":false,"explanation":"This incorrectly applies trends moving across a period. Down Groups 1 and 2, radius and shielding increase, weakening attraction to the outer electrons and increasing reactivity."},{"text":"

Atomic radius increases
Electron shielding decreases
Attraction between nucleus and outer electrons increases

","correct":false,"explanation":"A larger atom has more occupied shells, so shielding increases rather than decreases. This reduces nuclear attraction to the outer electrons and makes them easier to lose."},{"text":"

Atomic radius decreases
Electron shielding increases
Attraction between nucleus and outer electrons increases

","correct":false,"explanation":"Shielding does increase down the group, but atomic radius also increases and nuclear attraction to the outer electrons decreases. Consequently, the metals lose electrons more readily."}]},{"id":"886c743b-399b-4aa1-afb7-b3e2f049aaaf","cardType":"basic","points":3,"front":"Explain why potassium reacts more vigorously with water than sodium.","back":"

Potassium has a larger atomic radius and more shielding
It has a lower first ionisation energy
The outer electron is lost more easily

","mcqOptions":[{"text":"

Potassium has a larger atomic radius and more shielding
It has a lower first ionisation energy
The outer electron is lost more easily

","correct":true,"explanation":"Potassium’s extra electron shell increases radius and shielding. Its outer electron experiences weaker nuclear attraction than sodium’s, lowering ionisation energy and increasing reaction vigour."},{"text":"

Potassium has a smaller atomic radius and less shielding
It has a higher first ionisation energy
The outer electron is lost less easily

","correct":false,"explanation":"These trends are reversed. Potassium has an additional shell, giving greater radius and shielding, a lower first ionisation energy, and an outer electron that is lost more easily."},{"text":"

Potassium has a larger atomic radius and more shielding
It has a higher first ionisation energy
The outer electron is lost less easily

","correct":false,"explanation":"The structural comparison is right, but its consequence is reversed. Greater radius and shielding reduce nuclear attraction, giving potassium lower first ionisation energy and easier electron loss."},{"text":"

Potassium has a smaller atomic radius and more shielding
It has a lower first ionisation energy
The outer electron is gained more easily

","correct":false,"explanation":"Potassium’s radius is larger, not smaller, and Group 1 reactivity involves losing an outer electron rather than gaining one. Increased shielding makes that loss easier."}]},{"id":"19c28417-a241-45c2-ac8d-43185dfd333f","cardType":"basic","points":2,"front":"What are the products of the reaction between a Group 1 metal and water?","back":"A metal hydroxide and hydrogen gas.","mcqOptions":[{"text":"A metal hydroxide and hydrogen gas.","correct":true,"explanation":"A Group 1 metal displaces hydrogen from water, producing an alkaline solution of the soluble metal hydroxide and releasing hydrogen gas."},{"text":"A metal oxide and hydrogen gas.","correct":false,"explanation":"Metal oxides may form when Group 1 metals react with oxygen, but water produces a metal hydroxide. Hydrogen gas is released alongside it."},{"text":"A metal hydroxide and oxygen gas.","correct":false,"explanation":"The metal hydroxide is correctly identified, but the gas is not oxygen. The metal displaces hydrogen from water, so hydrogen gas is released."},{"text":"A metal hydride and oxygen gas.","correct":false,"explanation":"This confuses reactions with hydrogen and oxygen. A Group 1 metal reacting with water forms a metal hydroxide and releases hydrogen gas."}]},{"id":"2532fa39-66c6-4248-a0a8-7143c0acae30","cardType":"basic","points":1,"front":"Which Group 1 metal burns with a lilac flame during its reaction with water?","back":"Potassium","mcqOptions":[{"text":"Potassium","correct":true,"explanation":"Potassium produces the characteristic lilac flame during its vigorous reaction with water, providing a useful qualitative identification of the metal."},{"text":"Lithium","correct":false,"explanation":"Lithium is associated with a crimson-red flame rather than a lilac one. The Group 1 metal that burns with a lilac flame is potassium."},{"text":"Sodium","correct":false,"explanation":"Sodium produces a characteristic yellow-orange flame during its reaction with water. A lilac flame is the characteristic observation for potassium."},{"text":"Rubidium","correct":false,"explanation":"Rubidium reacts extremely violently and is associated with a red-violet flame, but the standard Group 1 metal identified by a lilac flame is potassium."}]},{"id":"a9966b29-7bb2-47a3-b22d-ee104989dcc4","cardType":"basic","points":1,"front":"State the formula of the oxide usually formed when lithium reacts with oxygen.","back":"$$\\mathrm{Li_2O}$","mcqOptions":[{"text":"$$\\mathrm{Li_2O}$","correct":true,"explanation":"Lithium reacts with oxygen to form lithium oxide, in which two lithium ions balance one oxide ion, giving the formula $\\mathrm{Li_2O}$."},{"text":"$$\\mathrm{Li_2O_2}$","correct":false,"explanation":"This is lithium peroxide, not the oxide usually produced. Lithium normally forms $\\mathrm{Li_2O}$ because two $\\mathrm{Li^+}$ ions are required for each $\\mathrm{O^{2-}}$ ion."},{"text":"$$\\mathrm{LiO_2}$","correct":false,"explanation":"This is the formula of lithium superoxide, whereas lithium is the Group 1 metal that usually forms the simple oxide $\\mathrm{Li_2O}$ when burned in oxygen."},{"text":"$$\\mathrm{LiO}$","correct":false,"explanation":"This incorrectly combines lithium and oxide ions in a 1:1 ratio. Charge balance requires two $\\mathrm{Li^+}$ ions per $\\mathrm{O^{2-}}$ ion, producing $\\mathrm{Li_2O}$."}]},{"id":"d5f7eb33-4a46-4101-be68-5671a3bc8760","cardType":"cloze","points":null,"front":"Magnesium reacts very slowly with cold water, but reacts with steam to form [ ] and hydrogen gas.","back":"Magnesium reacts very slowly with cold water, but reacts with steam to form magnesium oxide and hydrogen gas.","mcqOptions":[{"text":"magnesium oxide","correct":true,"explanation":"Steam oxidises magnesium to magnesium oxide while hydrogen ions from water are reduced to hydrogen gas: $\\mathrm{Mg + H_2O \\to MgO + H_2}$."},{"text":"magnesium hydroxide","correct":false,"explanation":"Magnesium hydroxide is associated with magnesium reacting slowly with liquid water. With steam, the greater heat favours formation of magnesium oxide and hydrogen gas."},{"text":"magnesium hydride","correct":false,"explanation":"Magnesium hydride is not formed by this metal-water reaction. Steam supplies oxygen to form magnesium oxide, while its hydrogen is released as hydrogen gas."},{"text":"magnesium peroxide","correct":false,"explanation":"Magnesium peroxide is not the normal product of magnesium reacting with steam. The balanced reaction produces magnesium oxide and hydrogen gas."}]},{"id":"75a934f5-7d74-4bc2-88ad-6bed912caa9d","cardType":"basic","points":2,"front":"Write the balanced chemical equation for the reaction of calcium with cold water.","back":"$$\\mathrm{Ca(s) + 2H_2O(l) \\to Ca(OH)_2(aq) + H_2(g)}$","mcqOptions":[{"text":"$$\\mathrm{Ca(s) + 2H_2O(l) \\to Ca(OH)_2(aq) + H_2(g)}$","correct":true,"explanation":"Calcium reacts with cold water to produce aqueous calcium hydroxide and hydrogen. Two water molecules provide the two hydroxide groups and the two hydrogen atoms in $\\mathrm{H_2}$."},{"text":"$$\\mathrm{Ca(s) + H_2O(l) \\to CaO(aq) + H_2(g)}$","correct":false,"explanation":"Calcium forms calcium hydroxide, not calcium oxide, in cold water. The correct equation is $\\mathrm{Ca(s) + 2H_2O(l) \\to Ca(OH)_2(aq) + H_2(g)}$."},{"text":"$$\\mathrm{Ca(s) + 2H_2O(l) \\to Ca(OH)_2(aq) + 2H_2(g)}$","correct":false,"explanation":"The coefficient of hydrogen has been doubled, leaving four hydrogen atoms on the left but six on the right. The balanced equation produces one $\\mathrm{H_2}$ molecule."},{"text":"$$\\mathrm{2Ca(s) + 2H_2O(l) \\to 2CaOH(aq) + H_2(g)}$","correct":false,"explanation":"Calcium forms $\\mathrm{Ca^{2+}}$, so the product is $\\mathrm{Ca(OH)_2}$ rather than calcium monohydroxide. The correct stoichiometric ratio is $1:2:1:1$."}]},{"id":"10b4b9a1-9baa-49d5-bcfd-adf5822dc1d9","cardType":"basic","points":1,"front":"Describe the appearance of magnesium burning in oxygen.","back":"It burns with a bright white flame.","mcqOptions":[{"text":"It burns with a bright white flame.","correct":true,"explanation":"Burning magnesium emits an intense bright white light because the oxidation reaction releases substantial energy, producing solid white magnesium oxide."},{"text":"It burns with a brick-red flame.","correct":false,"explanation":"A brick-red colour is associated with calcium compounds in flame tests, not burning magnesium. Magnesium burns with a bright white flame."},{"text":"It burns with a bright lilac flame.","correct":false,"explanation":"A lilac flame is characteristic of potassium compounds in a flame test. Burning magnesium instead produces a distinctive, intense bright white flame."},{"text":"It burns with a pale blue flame.","correct":false,"explanation":"A pale blue flame is associated with substances such as burning sulfur or carbon monoxide. Magnesium burning in oxygen gives a bright white flame."}]},{"id":"6f42699b-9559-404f-91f3-b071dafe324a","cardType":"basic","points":1,"front":"What products are formed when a Group 2 oxide reacts with an acid?","back":"A salt and water.","mcqOptions":[{"text":"A salt and water.","correct":true,"explanation":"Group 2 oxides are basic, so they neutralise acids. The oxide ions accept protons to form water, while the metal ions form a salt with the acid’s anions."},{"text":"A salt and hydrogen.","correct":false,"explanation":"Hydrogen is released when certain metals react with acids, not when basic metal oxides neutralise them. A Group 2 oxide and an acid produce a salt and water."},{"text":"A salt and carbon dioxide.","correct":false,"explanation":"Carbon dioxide is produced when a carbonate reacts with an acid. An oxide contains no carbonate ions, so its acid reaction produces a salt and water."},{"text":"A metal and water.","correct":false,"explanation":"An acid does not reduce a Group 2 oxide to the metal. This is an acid-base neutralisation reaction whose products are a salt and water."}]},{"id":"3262af7c-ab0f-4217-9823-22c8d7715f48","cardType":"basic","points":1,"front":"Why is magnesium hydroxide used in indigestion remedies?","back":"It neutralises excess stomach acid.","mcqOptions":[{"text":"It neutralises excess stomach acid.","correct":true,"explanation":"Magnesium hydroxide is a base that reacts with excess hydrochloric acid in the stomach, raising the pH and relieving symptoms caused by excess acidity."},{"text":"It prevents stomach acid production.","correct":false,"explanation":"Preventing acid secretion is the action of some medicines, such as proton-pump inhibitors, but not magnesium hydroxide. Magnesium hydroxide works by neutralising acid already present."},{"text":"It catalyses digestion of stomach proteins.","correct":false,"explanation":"Digestive proteases catalyse protein digestion; magnesium hydroxide does not. Its usefulness as an antacid comes from neutralising excess stomach acid through an acid-base reaction."},{"text":"It coats the stomach with an insoluble barrier.","correct":false,"explanation":"This confuses magnesium hydroxide with barrier-forming reflux treatments such as alginates. Magnesium hydroxide relieves indigestion by chemically neutralising excess stomach acid."}]},{"id":"10a1398c-f77d-49df-9e65-f5b9881682c2","cardType":"basic","points":2,"front":"State the products formed from the thermal decomposition of a Group 2 carbonate.","back":"Metal oxide and carbon dioxide.","mcqOptions":[{"text":"Metal oxide and carbon dioxide.","correct":true,"explanation":"Heating a Group 2 carbonate causes it to decompose into the corresponding metal oxide and carbon dioxide, following the general equation $\\mathrm{MCO_3 \\to MO + CO_2}$."},{"text":"Metal metal and carbon dioxide.","correct":false,"explanation":"Thermal decomposition does not reduce the Group 2 ion to its metal. The carbonate breaks down into the metal oxide and carbon dioxide."},{"text":"Metal oxide and carbon monoxide.","correct":false,"explanation":"Carbon remains at oxidation state +4 during this decomposition, so carbon dioxide rather than carbon monoxide forms. The other product is the corresponding metal oxide."},{"text":"Metal hydroxide and carbon dioxide.","correct":false,"explanation":"A hydroxide would require hydrogen, which is absent from a carbonate. Heating a Group 2 carbonate produces the metal oxide and carbon dioxide."}]},{"id":"c225611d-702f-4c23-8f0e-3e897769a474","cardType":"basic","points":3,"front":"What three products are formed when Group 2 nitrates are heated?","back":"

Metal oxide
Nitrogen dioxide
Oxygen

","mcqOptions":[{"text":"

Metal oxide
Nitrogen dioxide
Oxygen

","correct":true,"explanation":"Group 2 nitrates decompose on heating to the corresponding metal oxide, brown nitrogen dioxide gas and oxygen gas. This differs from the behaviour of most Group 1 nitrates."},{"text":"

Metal nitrite
Nitrogen dioxide
Oxygen

","correct":false,"explanation":"Formation of a metal nitrite is characteristic of most Group 1 nitrates, not Group 2 nitrates. Group 2 nitrates produce metal oxide, nitrogen dioxide and oxygen."},{"text":"

Metal oxide
Nitrogen
Oxygen

","correct":false,"explanation":"The nitrogen-containing gaseous product is brown nitrogen dioxide rather than nitrogen gas. Heating a Group 2 nitrate gives metal oxide, nitrogen dioxide and oxygen."},{"text":"

Metal oxide
Nitrogen dioxide
Hydrogen

","correct":false,"explanation":"No hydrogen is present in a nitrate, so hydrogen gas cannot form. The three products are the metal oxide, nitrogen dioxide and oxygen."}]},{"id":"053535a7-5756-4b40-9162-829913c8c1b5","cardType":"basic","points":2,"front":"Explain the trend in thermal stability of Group 2 carbonates down the group.","back":"

Stability increases
Cation size increases, so it polarises the anion less
The C-O bond is harder to break

","mcqOptions":[{"text":"

Stability increases
Cation size increases, so it polarises the anion less
The C-O bond is harder to break

","correct":true,"explanation":"Group 2 carbonate stability increases down the group because the larger cations have lower charge density and polarise the carbonate ion less, making its C-O bonds harder to break."},{"text":"

Stability decreases
Cation size increases, so it polarises the anion less
The C-O bond is harder to break

","correct":false,"explanation":"The stability trend has been reversed. Larger cations polarise the carbonate ion less strongly, so its C-O bonds are harder to break and thermal stability increases down Group 2."},{"text":"

Stability increases
Cation size increases, so it polarises the anion more
The C-O bond is harder to break

","correct":false,"explanation":"Increasing ionic radius lowers the cation's charge density, so its polarising power decreases rather than increases. Consequently, the C-O bonds are harder to break and stability increases."},{"text":"

Stability increases
Cation size increases, so it polarises the anion less
The C-O bond is easier to break

","correct":false,"explanation":"Reduced polarisation means the carbonate ion is distorted less, making its C-O bonds harder, not easier, to break. This causes thermal stability to increase down the group."}]},{"id":"be71087f-e7cc-4619-886d-509fd5ebf047","cardType":"cloze","points":null,"front":"In a flame test, [ ] and strontium ions both produce a crimson red colour.","back":"In a flame test, lithium and strontium ions both produce a crimson red colour.","mcqOptions":[{"text":"lithium","correct":true,"explanation":"Lithium ions produce a crimson red flame because their electrons emit characteristic visible wavelengths as they return from excited states to lower energy levels."},{"text":"calcium","correct":false,"explanation":"Calcium is tempting because its flame is red-toned, but calcium produces an orange-red or brick-red flame. Lithium is the ion that produces crimson red alongside strontium."},{"text":"potassium","correct":false,"explanation":"Potassium ions produce a lilac flame rather than a crimson red one. Lithium ions give the crimson red flame colour stated alongside strontium."},{"text":"sodium","correct":false,"explanation":"Sodium ions produce an intense yellow flame, which can mask other flame colours. Lithium, not sodium, produces a crimson red flame like strontium."}]},{"id":"648c2355-2f2f-43ee-a4c0-c5b450791729","cardType":"cloze","points":null,"front":"In a flame test, lithium and [ ] ions both produce a crimson red colour.","back":"In a flame test, lithium and strontium ions both produce a crimson red colour.","mcqOptions":[{"text":"strontium","correct":true,"explanation":"Strontium ions produce a crimson red flame because excited electrons release characteristic wavelengths of visible light when they return to lower energy levels."},{"text":"calcium","correct":false,"explanation":"Calcium can be confused with strontium because both give red-toned flames, but calcium is orange-red or brick-red. Strontium produces the crimson red colour."},{"text":"barium","correct":false,"explanation":"Barium ions give an apple-green flame rather than a crimson red flame. Strontium ions produce crimson red, matching the colour produced by lithium ions."},{"text":"sodium","correct":false,"explanation":"Sodium produces a characteristic intense yellow flame, not crimson red. Strontium is the Group 2 ion whose flame colour is crimson red like lithium's."}]},{"id":"b73edee3-7b0f-49fc-8688-21e80ffb8864","cardType":"basic","points":1,"front":"Identify the flame test colours for sodium and potassium ions.","back":"

Sodium: yellow
Potassium: lilac

","mcqOptions":[{"text":"

Sodium: yellow
Potassium: lilac

","correct":true,"explanation":"Sodium ions produce an intense yellow flame, while potassium ions produce a lilac flame. These colours arise from element-specific electronic transitions."},{"text":"

Sodium: lilac
Potassium: yellow

","correct":false,"explanation":"The two characteristic colours have been exchanged. Sodium gives an intense yellow flame, whereas potassium gives a lilac flame."},{"text":"

Sodium: yellow
Potassium: apple green

","correct":false,"explanation":"Sodium is correctly identified as yellow, but apple green is the flame colour of barium. Potassium produces a lilac flame."},{"text":"

Sodium: crimson red
Potassium: lilac

","correct":false,"explanation":"Potassium is correctly identified as lilac, but crimson red is associated with lithium and strontium. Sodium produces an intense yellow flame."}]},{"id":"736ff059-b4e3-428c-8f1b-67e71a64c6f8","cardType":"basic","points":1,"front":"What is the flame test colour for barium ions?","back":"Apple green","mcqOptions":[{"text":"Apple green","correct":true,"explanation":"Barium ions produce an apple-green flame as excited electrons return to lower energy levels and emit wavelengths characteristic of barium."},{"text":"Lilac","correct":false,"explanation":"Lilac is the characteristic flame colour of potassium ions, not barium ions. Barium ions produce an apple-green flame."},{"text":"Yellow","correct":false,"explanation":"Yellow is the characteristic flame colour of sodium ions and is often particularly intense. Barium ions instead produce an apple-green flame."},{"text":"Brick red","correct":false,"explanation":"Brick red is associated with calcium ions rather than barium ions. The characteristic flame-test colour for barium ions is apple green."}]},{"id":"078222f9-c314-4b84-97ce-698acd728104","cardType":"basic","points":1,"front":"Describe the solubility trend of Group 2 hydroxides down the group.","back":"Solubility increases.","mcqOptions":[{"text":"Solubility increases.","correct":true,"explanation":"Group 2 hydroxides become more soluble down the group because decreasing lattice enthalpy outweighs the accompanying decrease in the magnitude of hydration enthalpy."},{"text":"Solubility decreases.","correct":false,"explanation":"This reverses the hydroxide trend, possibly by confusing it with Group 2 sulfates. Hydroxide solubility increases because lattice enthalpy decreases sufficiently down the group."},{"text":"Solubility stays constant.","correct":false,"explanation":"Increasing cation radius changes the balance between lattice and hydration enthalpies, so solubility is not constant. Group 2 hydroxide solubility increases down the group."},{"text":"Solubility fluctuates.","correct":false,"explanation":"There is a consistent trend rather than irregular variation. Group 2 hydroxides become progressively more soluble down the group as their lattice enthalpies decrease."}]},{"id":"70944ed1-4673-4257-9360-1fce4915fe26","cardType":"basic","points":1,"front":"Describe the solubility trend of Group 2 sulfates down the group.","back":"Solubility decreases.","mcqOptions":[{"text":"Solubility decreases.","correct":true,"explanation":"Group 2 sulfate solubility decreases because hydration enthalpy becomes much less exothermic down the group, while lattice enthalpy changes relatively little due to the large sulfate ion."},{"text":"Solubility increases.","correct":false,"explanation":"This confuses the sulfate trend with that of Group 2 hydroxides. Sulfate solubility decreases as the cations' hydration enthalpies become less exothermic down the group."},{"text":"Solubility stays constant.","correct":false,"explanation":"Cation size affects hydration enthalpy, so sulfate solubility does not remain constant. It decreases down Group 2 as hydration becomes progressively less energetically favourable."},{"text":"Solubility fluctuates.","correct":false,"explanation":"The sulfate solubilities follow a clear downward trend rather than fluctuating. Decreasing hydration enthalpy makes the sulfates progressively less soluble down Group 2."}]},{"id":"814fe9de-6972-48d7-9c44-a16e68d6a87c","cardType":"basic","points":2,"front":"Write the ionic equation for the test for sulfate ions using barium ions.","back":"$$\\mathrm{Ba^{2+}(aq) + SO_4^{2-}(aq) \\to BaSO_4(s)}$","mcqOptions":[{"text":"$$\\mathrm{Ba^{2+}(aq) + SO_4^{2-}(aq) \\to BaSO_4(s)}$","correct":true,"explanation":"Barium ions and sulfate ions combine in a 1:1 ratio to form insoluble barium sulfate, observed as a white precipitate during the sulfate test."},{"text":"$$\\mathrm{Ba^{2+}(aq) + SO_4^{2-}(aq) \\to BaSO_4(aq)}$","correct":false,"explanation":"Barium sulfate is insoluble, so it must be shown as a solid precipitate rather than aqueous. The correct product is $\\mathrm{BaSO_4(s)}$."},{"text":"$$\\mathrm{Ba^{+}(aq) + SO_4^{2-}(aq) \\to BaSO_4(s)}$","correct":false,"explanation":"Barium forms $\\mathrm{Ba^{2+}}$, not $\\mathrm{Ba^+}$, and the displayed equation is not charge-balanced. Equal amounts of $\\mathrm{Ba^{2+}}$ and $\\mathrm{SO_4^{2-}}$ form $\\mathrm{BaSO_4(s)}$."},{"text":"$$\\mathrm{Ba^{2+}(aq) + SO_3^{2-}(aq) \\to BaSO_3(s)}$","correct":false,"explanation":"This equation involves sulfite ions, $\\mathrm{SO_3^{2-}}$, rather than the sulfate ions being tested. Sulfate reacts with barium ions to form $\\mathrm{BaSO_4(s)}$."}]},{"id":"2941d9d1-8fe1-4487-b027-d88928c08a3f","cardType":"cloze","points":null,"front":"Barium sulfate is safe for medical use because it is extremely [ ], preventing the release of toxic $\\mathrm{Ba^{2+}}$ ions.","back":"Barium sulfate is safe for medical use because it is extremely insoluble, preventing the release of toxic $\\mathrm{Ba^{2+}}$ ions.","mcqOptions":[{"text":"insoluble","correct":true,"explanation":"Barium sulfate has an exceptionally low solubility in water, so negligible quantities of toxic barium ions enter the body despite the compound being swallowed."},{"text":"soluble","correct":false,"explanation":"High solubility would allow barium sulfate to dissolve and release toxic ions. The correct term is insoluble because its very low solubility prevents significant barium ion absorption."},{"text":"bioavailable","correct":false,"explanation":"A bioavailable substance can be absorbed and used or distributed by the body, which would increase exposure to barium ions. Barium sulfate is safe because it is insoluble."},{"text":"volatile","correct":false,"explanation":"Volatility concerns how readily a substance vaporises, not whether it releases ions in the digestive system. Barium sulfate is safe because it is insoluble in water."}]},{"id":"fc9fd87c-01cd-428e-ab25-e73d3e1b3a0c","cardType":"cloze","points":null,"front":"Barium sulfate is safe for medical use because it is extremely insoluble, preventing the release of toxic [ ] ions.","back":"Barium sulfate is safe for medical use because it is extremely insoluble, preventing the release of toxic $\\mathrm{Ba^{2+}}$ ions.","mcqOptions":[{"text":"$$mathrm{Ba^{2+}}$","correct":true,"explanation":"Soluble barium compounds release $mathrm{Ba^{2+}}$ ions, which are toxic. Barium sulfate’s extremely low solubility prevents these ions from being released in significant quantities."},{"text":"$$mathrm{Ba^+}$","correct":false,"explanation":"Barium is a Group 2 element and therefore forms ions by losing two electrons, not one. The toxic ions whose release is prevented are $mathrm{Ba^{2+}}$ ions."},{"text":"$$mathrm{SO_4^{2-}}$","correct":false,"explanation":"Sulfate is the negative ion in barium sulfate, but it is not the toxic barium species of concern. Insolubility prevents the release of $mathrm{Ba^{2+}}$ ions."},{"text":"$$mathrm{Ca^{2+}}$","correct":false,"explanation":"Calcium ions are also Group 2 ions, but calcium is not the metal present in barium sulfate. The potentially toxic ions are $mathrm{Ba^{2+}}$ ions."}]}]}],"examCountdown":{"examDate":"2027-05-28","paperLabel":"Paper 1"},"heading":"The elements of Groups 1 and 2","title":"The elements of Groups 1 and 2 Flashcards","topbar":{"backHref":"/a-level/chemistry/edexcel/flashcards"},"topicSidebar":"$L38"}]}],"topicId":"2dd56d70-9f6c-5f0e-b542-b2f8b6bbb434","walled":"$L39"}],"$L3a","$L3b"]