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($e^-$).","correct":false,"explanation":"Electron transfer occurs in redox reactions, not in the Brønsted-Lowry definition of acids. A Brønsted-Lowry acid donates a proton, $H^+$."},{"text":"A hydrogen atom ($H$).","correct":false,"explanation":"A neutral hydrogen atom contains both a proton and an electron. A Brønsted-Lowry acid specifically donates a proton, represented by $H^+$."},{"text":"A hydroxide ion ($OH^-$).","correct":false,"explanation":"Hydroxide ions are commonly associated with alkaline solutions and proton acceptance. A Brønsted-Lowry acid donates a proton, $H^+$."}]},{"id":"4f36cce6-6b32-4e83-bd67-d2aff2471323","cardType":"basic","points":1,"front":"How many protons are donated by one molecule of a monobasic acid?","back":"One.","mcqOptions":[{"text":"One.","correct":true,"explanation":"A monobasic acid has one ionisable proton per molecule, so each molecule can donate exactly one proton in an acid-base reaction."},{"text":"Two.","correct":false,"explanation":"Donating two protons is characteristic of a dibasic acid, such as sulfuric acid. A monobasic acid donates only one proton per molecule."},{"text":"Three.","correct":false,"explanation":"Donating three protons is characteristic of a tribasic acid, such as phosphoric acid. A monobasic acid donates one proton per molecule."},{"text":"Zero.","correct":false,"explanation":"A substance donating no protons would not act as a Brønsted-Lowry acid. A monobasic acid can donate one proton from each molecule."}]},{"id":"9d8d4849-8fc5-4880-80bf-8e0a2cae9d34","cardType":"basic","points":1,"front":"What distinguishes a strong acid from a weak acid in terms of dissociation?","back":"A strong acid fully dissociates in aqueous solution.","mcqOptions":[{"text":"A strong acid fully dissociates in aqueous solution.","correct":true,"explanation":"A strong acid is essentially completely dissociated into ions in aqueous solution. A weak acid differs because only a proportion of its molecules dissociate."},{"text":"A strong acid partially dissociates in aqueous solution.","correct":false,"explanation":"Partial dissociation is the defining behaviour of a weak acid. A strong acid fully dissociates into ions in aqueous solution."},{"text":"A strong acid does not dissociate in aqueous solution.","correct":false,"explanation":"Failure to dissociate would prevent the substance from producing hydrogen ions in solution. A strong acid instead fully dissociates in aqueous solution."},{"text":"A strong acid fully dissolves in aqueous solution.","correct":false,"explanation":"Dissolving describes solubility and is not the same as dissociation into ions. A strong acid is distinguished by fully dissociating in aqueous solution."}]},{"id":"438c599a-3c9f-4420-b68e-52bbe457bfa9","cardType":"basic","points":1,"front":"For a strong monobasic acid, how does the hydrogen ion concentration $[H^+]$ relate to the initial acid concentration?","back":"They are equal ($[H^+] = [Acid]$).","mcqOptions":[{"text":"They are equal ($[H^+] = [Acid]$).","correct":true,"explanation":"A strong monobasic acid fully dissociates and releases one $H^+$ ion per acid molecule, so $[H^+]$ equals the initial acid concentration."},{"text":"The hydrogen ion concentration is double the acid concentration ($[H^+] = 2[Acid]$).","correct":false,"explanation":"This incorrectly treats the acid as dibasic, releasing two protons per molecule. A strong monobasic acid releases one, so $[H^+] = [Acid]$."},{"text":"The hydrogen ion concentration is half the acid concentration ($[H^+] = \\frac{1}{2}[Acid]$).","correct":false,"explanation":"Halving assumes only half the acid dissociates, which describes partial dissociation. A strong monobasic acid fully dissociates, making $[H^+] = [Acid]$."},{"text":"The hydrogen ion concentration is ten times the acid concentration ($[H^+] = 10[Acid]$).","correct":false,"explanation":"A factor of ten is relevant to changes on the logarithmic pH scale, not acid dissociation stoichiometry. Here, $[H^+] = [Acid]$."}]},{"id":"0cd4e513-350d-40f2-acb6-25dd91336fe9","cardType":"cloze","points":null,"front":"To convert volume from $\\text{cm}^3$ to $\\text{dm}^3$ for concentration calculations, you must [ ].","back":"To convert volume from $\\text{cm}^3$ to $\\text{dm}^3$ for concentration calculations, you must divide by 1000.","mcqOptions":[{"text":"divide by 1000","correct":true,"explanation":"One cubic decimetre contains $1000\\text{ cm}^3$, so converting a volume from $\\text{cm}^3$ to $\\text{dm}^3$ requires division by 1000."},{"text":"multiply by 1000","correct":false,"explanation":"Multiplying by 1000 converts $\\text{dm}^3$ to $\\text{cm}^3$, which is the reverse conversion. From $\\text{cm}^3$ to $\\text{dm}^3$, divide by 1000."},{"text":"divide by 100","correct":false,"explanation":"This treats the conversion as if only two factors of ten were involved. Since volume is cubic, $1000\\text{ cm}^3 = 1\\text{ dm}^3$, so divide by 1000."},{"text":"divide by 1,000,000","correct":false,"explanation":"Division by 1,000,000 applies when converting $\\text{cm}^3$ to $\\text{m}^3$. To obtain $\\text{dm}^3$, divide the volume by 1000."}]},{"id":"8d9a7b72-10a6-462f-a25e-deb457e706d6","cardType":"cloze","points":null,"front":"In pH calculations, acid concentration must be measured in the units $\\text{[...]}$.","back":"In pH calculations, acid concentration must be measured in the units $\\boxed{\\text{mol dm}^{-3}}$.","mcqOptions":[{"text":"$$\\text{mol dm}^{-3}$","correct":true,"explanation":"Hydrogen ion concentration in the pH equation is expressed as an amount concentration in moles per cubic decimetre, written $\\text{mol dm}^{-3}$."},{"text":"$$\\text{mol cm}^{-3}$","correct":false,"explanation":"Using cubic centimetres would make the numerical concentration 1000 times smaller than the standard value. pH calculations use concentration in $\\text{mol dm}^{-3}$."},{"text":"$$\\text{g dm}^{-3}$","correct":false,"explanation":"Grams per cubic decimetre measure mass concentration, not amount concentration. The pH equation requires hydrogen ion concentration in $\\text{mol dm}^{-3}$."},{"text":"$$\\text{mol m}^{-3}$","correct":false,"explanation":"Moles per cubic metre are valid SI concentration units but give a numerical value 1000 times larger. Standard pH calculations use $\\text{mol dm}^{-3}$."}]},{"id":"77bb7038-c1a1-4adf-b135-88556edf1a96","cardType":"basic","points":1,"front":"State the mathematical equation used to define pH.","back":"$$\\text{pH} = -\\log_{10}[H^+]$","mcqOptions":[{"text":"$$\\text{pH} = -\\log_{10}[H^+]$","correct":true,"explanation":"pH is defined as the negative base-10 logarithm of the hydrogen ion concentration, with $[H^+]$ measured in $\\text{mol dm}^{-3}$."},{"text":"$$\\text{pH} = \\log_{10}[H^+]$","correct":false,"explanation":"Omitting the negative sign would give negative pH values for ordinary dilute acids. The definition is $\\text{pH} = -\\log_{10}[H^+]$."},{"text":"$$\\text{pH} = -\\log_{10}[OH^-]$","correct":false,"explanation":"The negative logarithm of hydroxide ion concentration defines pOH, not pH. pH is calculated using $\\text{pH} = -\\log_{10}[H^+]$."},{"text":"$$\\text{pH} = 10^{-[H^+]}$","correct":false,"explanation":"This confuses the pH definition with its inverse exponential relationship. pH is found by taking the negative base-10 logarithm: $\\text{pH} = -\\log_{10}[H^+]$."}]},{"id":"c9483397-9646-4cea-8002-aaed3a44efb9","cardType":"cloze","points":null,"front":"To calculate hydrogen ion concentration from a known pH, the rearranged equation is $\\text{[...]}$.","back":"To calculate hydrogen ion concentration from a known pH, the rearranged equation is $\\boxed{[H^+] = 10^{-\\text{pH}}}$.","mcqOptions":[{"text":"$$[H^+] = 10^{-\\text{pH}}$","correct":true,"explanation":"Reversing the base-10 logarithm in the pH definition gives $[H^+] = 10^{-\\text{pH}}$, with the concentration expressed in $\\text{mol dm}^{-3}$."},{"text":"$$[H^+] = 10^{\\text{pH}}$","correct":false,"explanation":"This loses the negative sign from the pH definition and produces unrealistically large concentrations. The correct rearrangement is $[H^+] = 10^{-\\text{pH}}$."},{"text":"$$[H^+] = -\\log_{10}(\\text{pH})$","correct":false,"explanation":"This applies another logarithm instead of using the inverse operation. Exponentiating base 10 gives the correct rearrangement, $[H^+] = 10^{-\\text{pH}}$."},{"text":"$$[H^+] = 10^{-\\text{pOH}}$","correct":false,"explanation":"Using pOH gives the hydroxide ion concentration rather than the hydrogen ion concentration. From pH, the correct relationship is $[H^+] = 10^{-\\text{pH}}$."}]},{"id":"c8fab931-7b7a-4b60-8b3d-94f46ce69e5c","cardType":"basic","points":1,"front":"If the pH of a solution decreases by one unit, what happens to $[H^+]$?","back":"It increases by a factor of 10.","mcqOptions":[{"text":"It increases by a factor of 10.","correct":true,"explanation":"Each decrease of one pH unit represents a tenfold increase in hydrogen ion concentration because pH is defined as minus the base-10 logarithm of $[H^+]$."},{"text":"It decreases by a factor of 10.","correct":false,"explanation":"A pH decrease means greater acidity, not lower hydrogen ion concentration. The logarithmic scale shows that one unit corresponds specifically to a tenfold increase."},{"text":"It increases by a factor of 100.","correct":false,"explanation":"A ten-unit pH change would produce a factor of $10^{10}$, not a one-unit change. One pH unit corresponds to a factor of 10."},{"text":"It remains unchanged.","correct":false,"explanation":"Because pH is logarithmic, changing it by one unit changes $[H^+]$ tenfold. The concentration therefore cannot remain unchanged."}]},{"id":"aac8ecbe-9517-47cb-b238-ad78f8e5db83","cardType":"basic","points":2,"front":"Compare the acidity of solution A (pH 2.00) and solution B (pH 5.00).","back":"Solution A has 1000 times higher $[H^+]$ than solution B.","mcqOptions":[{"text":"Solution A has 1000 times higher $[H^+]$ than solution B.","correct":true,"explanation":"The pH difference is three units, and each unit represents a tenfold change. Therefore solution A has $10^3$, or 1000, times the hydrogen ion concentration."},{"text":"Solution A has 100 times higher $[H^+]$ than solution B.","correct":false,"explanation":"A difference of two pH units would give a factor of 100. Here the difference is three units, so solution A is 1000 times higher in $[H^+]$."},{"text":"Solution A has 3 times higher $[H^+]$ than solution B.","correct":false,"explanation":"The pH difference is three units, not threefold concentration change. Each unit multiplies $[H^+]$ by 10, giving a factor of 1000."},{"text":"Solution B has 1000 times higher $[H^+]$ than solution A.","correct":false,"explanation":"The lower pH solution is more acidic and has the greater hydrogen ion concentration. Reversing the comparison would make solution B appear incorrectly more acidic."}]},{"id":"d904f0ed-4c79-48da-8f20-b191573a4c20","cardType":"cloze","points":null,"front":"When calculating pH, use the $\\text{[...]}$ button on a calculator rather than the $\\ln$ button.","back":"When calculating pH, use the $\\boxed{\\log_{10}}$ button on a calculator rather than the $\\ln$ button.","mcqOptions":[{"text":"$$\\log_{10}$","correct":true,"explanation":"The pH definition uses the common logarithm, which is the base-10 logarithm. Selecting this button gives the required value for pH calculations."},{"text":"$$\\ln$","correct":false,"explanation":"The natural logarithm uses base $e$, whereas pH is defined using base 10. Using it produces a different numerical result unless a conversion is made."},{"text":"$$\\log_{2}$","correct":false,"explanation":"A base-2 logarithm is not used in the definition of pH. The required logarithm is the common logarithm with base 10."},{"text":"$$10^x$","correct":false,"explanation":"An inverse logarithm converts a logarithmic value back to a concentration; it does not calculate pH from hydrogen ion concentration."}]},{"id":"fe325e90-5701-48f2-8776-2aaf515ad44b","cardType":"cloze","points":null,"front":"When calculating pH, use the $\\log_{10}$ button on a calculator rather than the $\\text{[...]}$ button.","back":"When calculating pH, use the $\\log_{10}$ button on a calculator rather than the $\\boxed{\\ln}$ button.","mcqOptions":[{"text":"$$\\ln$","correct":true,"explanation":"The natural logarithm uses base $e$, but pH calculations require the common logarithm with base 10. Therefore the natural-logarithm button fills the blank."},{"text":"$$\\log_{10}$","correct":false,"explanation":"Using the base-10 logarithm in both parts would make the instruction say to use the same button rather than contrasting it with the correct alternative."},{"text":"$$\\log_{2}$","correct":false,"explanation":"The base-2 logarithm is unrelated to the standard definition of pH. The relevant alternative button here is the natural logarithm, $\\ln$."},{"text":"$$e^x$","correct":false,"explanation":"An exponential or inverse-log function is not the natural logarithm. The incorrect calculator button being contrasted with $\\log_{10}$ is $\\ln$."}]},{"id":"c98d755c-e6e9-4914-939d-5aa28fab11bb","cardType":"basic","points":2,"front":"Contrast the terms strong and concentrated when describing an acid.","back":"

Strong: The extent of dissociation into ions (complete).
Concentrated: The amount of acid moles dissolved per unit volume.

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

Strong: The extent of dissociation into ions (complete).
Concentrated: The amount of acid moles dissolved per unit volume.

","correct":true,"explanation":"Strong describes complete dissociation into ions, while concentrated describes how many moles of acid are present per unit volume. These are different properties."},{"text":"

Strong: The amount of acid moles dissolved per unit volume.
Concentrated: The extent of dissociation into ions (complete).

","correct":false,"explanation":"This option swaps the definitions: amount per volume describes concentration, while complete ionisation describes strength. Strength and concentration must not be treated as interchangeable."},{"text":"

Strong: A high number of acid moles per unit volume.
Concentrated: Complete dissociation into ions.

","correct":false,"explanation":"A high amount per unit volume defines concentration, not strength. Concentrated acids may be weak if they only partially dissociate into ions."},{"text":"

Strong: An acid that always has a lower pH.
Concentrated: The fraction dissociated into ions.

","correct":false,"explanation":"pH depends on concentration as well as acid strength, so strength is not simply an acid having lower pH. Concentration is not the fraction dissociated."}]},{"id":"49392f2c-fbeb-495b-bc89-9abc397958b3","cardType":"basic","points":1,"front":"Why is the pH of pure water exactly 7.00 only at 298 K?","back":"The dissociation of water (and thus pH) is temperature-dependent.","mcqOptions":[{"text":"The dissociation of water (and thus pH) is temperature-dependent.","correct":true,"explanation":"Water ionises to different extents at different temperatures, changing its neutral $[H^+]$ and therefore its neutral pH. pH 7.00 applies to pure water at 298 K."},{"text":"The dissociation of water (and thus pH) is temperature-independent.","correct":false,"explanation":"Water dissociation is not temperature-independent. The ionisation constant and neutral pH vary with temperature, so pH 7.00 is not universal."},{"text":"The dissociation of water (and thus pH) is fixed at 273 K.","correct":false,"explanation":"273 K is the freezing point of water, not the reference temperature for pH 7.00. The stated reference temperature is 298 K."},{"text":"The dissociation of water (and thus pH) is always exactly 7.00.","correct":false,"explanation":"Pure water remains neutral at other temperatures, but neutral does not necessarily mean pH 7.00 because the ion concentrations change with temperature."}]},{"id":"fa7837ed-26f9-4f52-b566-4ebcce9996dd","cardType":"basic","points":2,"front":"Calculate the pH of $0.050\\ \\text{mol dm}^{-3}$ nitric acid ($HNO_3$).","back":"1.30 (as $[H^+] = 0.050$).","mcqOptions":[{"text":"1.30 (as $[H^+] = 0.050$).","correct":true,"explanation":"Nitric acid is strong and monobasic, so it fully dissociates and $[H^+] = 0.050$. Applying $\\mathrm{pH}=-\\log_{10}[H^+]$ gives 1.30."},{"text":"1.00 (as $[H^+] = 0.100$).","correct":false,"explanation":"A pH of 1.00 corresponds to $[H^+] = 0.100$. The acid concentration is 0.050, so the logarithm gives 1.30 instead."},{"text":"2.60 (as $[H^+] = 0.0025$).","correct":false,"explanation":"Squaring 0.050 gives 0.0025 and leads to 2.60, but nitric acid dissociates once without squaring its concentration. The correct pH is 1.30."},{"text":"0.050 (as $[H^+] = 1.30$).","correct":false,"explanation":"The concentration is the hydrogen ion concentration for this strong monobasic acid, but it must be transformed using the negative logarithm. The resulting pH is 1.30."}]},{"id":"aaa09471-8a42-43a7-a385-8f3f4249a8de","cardType":"basic","points":2,"front":"A strong monobasic acid has a pH of 1.70. Calculate its concentration.","back":"0.020 $\\text{mol dm}^{-3}$ ($10^{-1.70}$).","mcqOptions":[{"text":"0.020 $\\text{mol dm}^{-3}$ ($10^{-1.70}$).","correct":true,"explanation":"A strong monobasic acid fully dissociates, so its concentration equals $[H^+]$. Using $[H^+] = 10^{-\\mathrm{pH}}$ gives $10^{-1.70}=0.020\\ \\text{mol dm}^{-3}$."},{"text":"0.0050 $\\text{mol dm}^{-3}$ ($10^{-2.30}$).","correct":false,"explanation":"The value 0.0050 corresponds approximately to $10^{-2.30}$, effectively using a pH that is too high. For pH 1.70, the concentration is 0.020."},{"text":"0.050 $\\text{mol dm}^{-3}$ ($10^{-1.30}$).","correct":false,"explanation":"This uses $10^{-1.30}$ rather than $10^{-1.70}$, a likely logarithm or pH transcription error. The correct concentration is $10^{-1.70}$, equal to 0.020."},{"text":"50 $\\text{mol dm}^{-3}$ ($10^{1.70}$).","correct":false,"explanation":"Using a positive exponent reverses the relationship between pH and hydrogen ion concentration. The concentration must be $10^{-1.70}$, not $10^{1.70}$."}]},{"id":"c6ba2b51-dabb-4892-962b-6f22f7ddaa64","cardType":"basic","points":1,"front":"Which formula is used to find the new concentration when a strong acid is diluted with water?","back":"$$c_1V_1 = c_2V_2$","mcqOptions":[{"text":"$$c_1V_1 = c_2V_2$","correct":true,"explanation":"Dilution conserves the number of moles of acid, so initial concentration multiplied by initial volume equals final concentration multiplied by final volume."},{"text":"$$c_1V_2 = c_2V_1$","correct":false,"explanation":"This incorrectly swaps the initial and final volumes. Conservation of moles during dilution gives $c_1V_1 = c_2V_2$."},{"text":"$$\\frac{c_1}{V_1} = \\frac{c_2}{V_2}$","correct":false,"explanation":"Dividing concentration by volume does not represent the number of moles present. Multiplication is required, giving $c_1V_1 = c_2V_2$."},{"text":"$$c_1 + V_1 = c_2 + V_2$","correct":false,"explanation":"Concentration and volume cannot be added because they have different units. Dilution conserves moles, so the appropriate relationship is $c_1V_1 = c_2V_2$."}]},{"id":"7ab30cf7-a08e-4060-abe4-01bf93073a30","cardType":"basic","points":1,"front":"What is the effect on pH when a strong acid is diluted by a factor of 10?","back":"The pH increases by 1 unit.","mcqOptions":[{"text":"The pH increases by 1 unit.","correct":true,"explanation":"A tenfold dilution divides $[H^+]$ by 10. Because pH is $-\\log_{10}[H^+]$, reducing the concentration by one power of ten raises pH by 1."},{"text":"The pH decreases by 1 unit.","correct":false,"explanation":"This reverses the effect of dilution: adding water lowers $[H^+]$ rather than raising it. A tenfold decrease in $[H^+]$ increases the pH by 1 unit."},{"text":"The pH increases by 10 units.","correct":false,"explanation":"This treats pH as directly proportional to concentration, overlooking its logarithmic definition. A tenfold dilution changes pH by 1 unit, not 10 units."},{"text":"The pH remains unchanged.","correct":false,"explanation":"Dilution lowers the hydrogen ion concentration, so the pH cannot remain constant. A tenfold reduction in $[H^+]$ increases the pH by 1 unit."}]},{"id":"1895f017-72cb-4260-968a-7842c5a78b3e","cardType":"basic","points":2,"front":"Why does the assumption $[H^+] = [Acid]$ fail for extremely dilute strong acids (e.g., $10^{-8}\\ \\text{mol dm}^{-3}$)?","back":"The $H^+$ ions provided by the dissociation of water become significant.","mcqOptions":[{"text":"The $H^+$ ions provided by the dissociation of water become significant.","correct":true,"explanation":"Water establishes its own equilibrium concentration of hydrogen ions. When the acid concentration is around or below $10^{-7}\\ \\text{mol dm}^{-3}$, this contribution can no longer be neglected."},{"text":"The $H^+$ ions retained by the undissociated acid molecules become significant.","correct":false,"explanation":"Strong acids remain essentially fully dissociated at very low concentrations, so undissociated acid is not the cause. The hydrogen ions arising from water become significant instead."},{"text":"The $H^+$ ions provided by the dissociation of hydroxide ions become significant.","correct":false,"explanation":"Hydroxide ions do not supply the additional hydrogen ions omitted by the assumption. The relevant contribution is the $H^+$ produced by water’s own dissociation equilibrium."},{"text":"The $H^+$ ions lost through the evaporation of the acid become significant.","correct":false,"explanation":"The failure is not caused by acid molecules evaporating from a dilute aqueous solution. It occurs because water’s dissociation supplies a significant proportion of the total $H^+$."}]},{"id":"9abef474-85f6-4714-8d2b-c136d2f0c0f0","cardType":"basic","points":1,"front":"Is it possible for a solution to have a pH value less than 0?","back":"Yes, if the acid is very concentrated ($[H^+] > 1\\ \\text{mol dm}^{-3}$).","mcqOptions":[{"text":"Yes, if the acid is very concentrated ($[H^+] > 1\\ \\text{mol dm}^{-3}$).","correct":true,"explanation":"The pH scale has no absolute lower boundary at zero. If $[H^+]$ exceeds $1\\ \\text{mol dm}^{-3}$, then $-\\log_{10}[H^+]$ is negative."},{"text":"No, because the pH scale begins at 0.","correct":false,"explanation":"Zero is not a fixed lower limit of the pH scale. A sufficiently concentrated acid with $[H^+] > 1\\ \\text{mol dm}^{-3}$ can have a negative pH."},{"text":"Yes, if the acid is very dilute ($[H^+] < 1\\ \\text{mol dm}^{-3}$).","correct":false,"explanation":"A hydrogen ion concentration below $1\\ \\text{mol dm}^{-3}$ gives a positive pH, not a negative one. Negative pH requires $[H^+] > 1\\ \\text{mol dm}^{-3}$."},{"text":"Yes, if the acid is fully dissociated ($[H^+] < 1\\ \\text{mol dm}^{-3}$).","correct":false,"explanation":"Complete dissociation alone does not guarantee a negative pH; ordinary dilute strong acids still have positive pH values. The concentration must make $[H^+]$ exceed $1\\ \\text{mol dm}^{-3}$."}]},{"id":"f428e783-5b62-4c79-9340-7020bc50e8e4","cardType":"cloze","points":null,"front":"For quantitative work, a pH meter must be [ ] using buffer solutions.","back":"For quantitative work, a pH meter must be calibrated using buffer solutions.","mcqOptions":[{"text":"calibrated","correct":true,"explanation":"Calibration compares the meter’s readings with solutions of accurately known pH, allowing its response to be adjusted before quantitative measurements are made."},{"text":"neutralised","correct":false,"explanation":"Neutralising a meter is not a valid preparation procedure and could damage or contaminate the electrode. It must be calibrated against buffer solutions of known pH."},{"text":"titrated","correct":false,"explanation":"Titration determines an unknown concentration by reacting substances in measured amounts; it does not prepare the instrument. A pH meter must instead be calibrated using buffers."},{"text":"diluted","correct":false,"explanation":"Dilution changes a solution’s concentration and is not an adjustment performed on a measuring instrument. The meter must be calibrated with buffers of known pH."}]},{"id":"3703473d-f8da-4a64-932a-699d7ac80cf6","cardType":"cloze","points":null,"front":"For quantitative work, a pH meter must be calibrated using [ ].","back":"For quantitative work, a pH meter must be calibrated using buffer solutions.","mcqOptions":[{"text":"buffer solutions","correct":true,"explanation":"Buffer solutions have stable, accurately known pH values, so they provide reliable reference points for adjusting and checking a pH meter."},{"text":"distilled water","correct":false,"explanation":"Distilled water does not provide several stable reference pH values and can absorb carbon dioxide from air. Calibration requires buffer solutions with accurately known pH values."},{"text":"strong acid solutions","correct":false,"explanation":"Strong acid solutions are not the standard stable references used across the meter’s working range. Calibration uses buffer solutions with accurately known pH values."},{"text":"indicator solutions","correct":false,"explanation":"Indicators estimate pH through colour changes rather than supplying precise reference values for an electrode. A pH meter is calibrated using buffer solutions."}]}]}],"examCountdown":{"examDate":"2027-05-11","paperLabel":"Paper 1"},"heading":"pH and strong acids","title":"pH and strong acids Flashcards","topbar":{"backHref":"/as-level/chemistry/edexcel/flashcards"},"topicSidebar":"$L37"}]}],"topicId":"42d11e08-213a-5516-afca-e4014ceb7d2f","walled":"$L38"}],"$L39","$L3a"]