Acids Bases and pH — Free Chemistry Review Games.

A neutral solution has a pH of 7, where the concentrations of H+ and OH- ions are equal.

Acidic solutions have a pH below 7, with lower numbers indicating stronger acidity.

Acids release hydrogen ions (H+) when dissolved in water.

When an acid reacts with a base, they neutralize each other to produce a salt and water.

Baking soda (sodium bicarbonate) is a common base with a pH above 7.

The pH scale measures the concentration of H+ ions; pH = -log[H+].

Bases produce hydroxide ions (OH-) in water or accept hydrogen ions.

A strong acid completely dissociates into ions in water, like HCl, HNO3, and H2SO4.

Each pH unit represents a tenfold change in H+ concentration, so pH 3 is 10^2 = 100 times more acidic than pH 5.

An indicator is a substance (like litmus or phenolphthalein) that changes color at different pH values to show whether a solution is acidic or basic.

The Bronsted-Lowry definition states that an acid is a proton donor, giving H+ to another molecule.

A buffer contains a weak acid and its conjugate base (or vice versa) and maintains a relatively stable pH.

pH = -log[H+] = -log(10^-4) = 4.

A conjugate acid-base pair consists of two species that differ by exactly one proton; the acid donates it, and the conjugate base is what remains.

Weak acids establish an equilibrium between ionized and un-ionized forms, with most molecules remaining undissociated.

The Arrhenius theory defines a base as a substance that increases the concentration of hydroxide ions (\(OH^-\)) when dissolved in water. "Hydrogen ions (\(H^+\))" is wrong because that describes an Arrhenius acid, not a base. Remember that Arrhenius theory is specific to aqueous solutions and defines acids and bases by which ion they release.

Neutralization is the specific term for the reaction in which an acid and a base combine to form a salt and water, driven by \(H^+\) and \(OH^-\) combining to form \(H_2O\). "Oxidation" is wrong because that involves electron transfer and a change in oxidation state, not proton transfer. This reaction type is central to titration and stoichiometry problems on the exam.

A pH of 14 represents the highest hydroxide ion concentration relative to hydrogen ion concentration on the standard 0-14 scale, making it the most basic value listed. "0" is wrong because it represents the most acidic end of the scale, not the most basic. Students should remember the scale runs from highly acidic (0) to highly basic (14), with 7 as neutral.

Hydrochloric acid ($HCl$) is a strong acid that fully dissociates in water to release \(H^+\) ions, making it a common acidic reagent in labs. "Sodium hydroxide ($NaOH$)" is wrong because it is a strong base that releases \(OH^-\) ions instead. Recognizing common acids and bases by formula helps quickly classify substances on the exam.

Litmus paper turns red in the presence of an acidic solution because the dye molecule changes structure in response to excess \(H^+\) ions. "Blue" is wrong because blue litmus indicates a basic, not acidic, environment. Indicators like litmus are simple qualitative tools for classifying solutions before precise pH measurement.

A pH greater than 7 indicates a lower concentration of \(H^+\) relative to \(OH^-\), which defines a basic or alkaline solution on the standard scale. "Equal to 7" is wrong because that value represents a neutral solution, such as pure water, not a basic one. Students should memorize that the neutral point is 7, with values above being basic and below being acidic.

The Bronsted-Lowry theory defines acids and bases in terms of proton (\(H^+\)) transfer, where the acid donates a proton and the base accepts one. "Electron" is wrong because electron transfer characterizes redox reactions, not the Bronsted-Lowry acid-base framework. This proton-transfer concept broadens acid-base chemistry beyond the Arrhenius definition, which only applies to aqueous solutions.

Pure water at 25 degrees Celsius has equal concentrations of \(H^+\) and \(OH^-\) ions, both at \(1\times10^{-7}\) M, giving it a pH of 7. "14" is wrong because that value corresponds to a strongly basic solution with a much higher \(OH^-\) concentration than pure water. Pure water serves as the reference point for neutrality on the pH scale.

An amphoteric substance, such as water, can either donate or accept a proton depending on what it reacts with, allowing it to behave as an acid or a base. "A strong acid" is wrong because a strong acid only donates protons and does not act as a base. Water's amphoteric nature is a key reason it plays a central role in acid-base chemistry.

When an acid reacts with a metal hydroxide, the \(H^+\) from the acid combines with \(OH^-\) from the base to form water, while the remaining ions form a salt. "Only water" is wrong because it ignores the salt that forms from the leftover cation and anion. This salt-plus-water pattern is the defining outcome of a classic neutralization reaction.

A weak base only partially ionizes in water, establishing an equilibrium between the unionized base and its ionized form rather than fully dissociating. "A base with a pH of exactly 14" is wrong because pH depends on concentration and ionization degree together, and no base is defined solely by that specific value. Recognizing partial versus complete ionization is essential for distinguishing strong from weak acids and bases.

A pH meter uses a specialized electrode sensitive to hydrogen ion activity to give a precise, quantitative pH reading of a solution. "A thermometer" is wrong because it measures temperature, not hydrogen ion concentration. While indicators give a rough estimate, a pH meter provides more accurate numerical results in laboratory settings.

Since \(pH = -\log[H^+]\), taking the negative log of \(1\times10^{-2}\) gives a pH of 2. "12" is wrong because it would correspond to a hydroxide concentration calculation, not this hydrogen ion concentration directly. Students must remember the formula \(pH = -\log[H^+]\) to convert between concentration and pH quickly.

Using \(M_aV_a = M_bV_b\), moles of base equal \(0.2\times0.05 = 0.01\) mol, so the acid volume needed is \(0.01/0.1 = 0.1\) L, or 100 mL. "50 mL" is wrong because it would only supply \(0.1\times0.05=0.005\) mol of acid, insufficient to neutralize the \(0.01\) mol of base present. This stoichiometric relationship underlies all acid-base titration calculations.

Since \(pH=-\log[H^+]\), an increase in \([H^+]\) produces a smaller (more negative before the sign flip) logarithm, which numerically lowers the pH value. "Because the pH scale increases as acidity increases" is wrong because the scale actually decreases as acidity, and thus \([H^+]\), increases. This inverse logarithmic relationship is essential for correctly interpreting pH changes in any acid-base scenario.

Ammonia acts as a Bronsted-Lowry base because its lone pair of electrons on nitrogen accepts a proton from water, forming \(NH_4^+\) and \(OH^-\). "It releases \(OH^-\) ions directly into solution" is wrong because ammonia does not contain \(OH^-\) in its structure; the hydroxide is generated indirectly through the proton-transfer reaction with water. This distinguishes Bronsted-Lowry bases, which need not contain hydroxide groups, from Arrhenius bases.

Each whole unit decrease in pH corresponds to a tenfold increase in \([H^+]\) because the pH scale is logarithmic in base 10, so a change from pH 5 to pH 4 means the concentration increased by a factor of 10. "100 times" is wrong because that would require a two-unit pH change, not one. Students should always convert pH differences into powers of ten to find concentration ratios.

Strength describes how completely an acid ionizes in water, whereas concentration describes the quantity of acid dissolved in a given volume, and these are independent properties. "Strong and concentrated mean the same thing in chemistry" is wrong because a dilute solution of a strong acid can have low \([H^+]\) while a concentrated weak acid may still have moderate \([H^+]\). This distinction is a common point of confusion that students must master to correctly compare acid solutions.

A buffer resists large pH changes because its conjugate base component reacts with and neutralizes the added \(H^+\) ions, keeping the pH relatively stable. "The pH drops dramatically, similar to adding acid to pure water" is wrong because pure water lacks the conjugate acid-base pair needed to absorb the added protons. This resistance to pH change is the defining functional property of buffer systems used throughout chemistry and biology.

Acetic acid and sodium acetate form a buffer because they provide a weak acid and its conjugate base in solution, allowing the system to neutralize added acid or base. "Hydrochloric acid and sodium hydroxide" is wrong because these are a strong acid and strong base that fully react to completion rather than establishing an equilibrium buffer pair. Effective buffers always pair a weak acid or base with its conjugate partner.

$HCl$ acts as the Bronsted-Lowry acid because it donates a proton (\(H^+\)) to the hydroxide ion during the reaction. "$NaOH$" is wrong because it acts as the base by accepting the proton donated by $HCl$. Identifying the proton donor versus acceptor in a written reaction is a key skill for classifying acid-base mechanisms.

Since a strong acid ionizes almost completely, its conjugate base has very little affinity for recapturing a proton, making it a weak base by comparison. "Because conjugate bases only form from weak acids" is wrong because every acid, strong or weak, produces a conjugate base upon losing a proton. This inverse relationship between acid strength and conjugate base strength is central to predicting the direction of acid-base equilibria.

Since $pH + pOH = 14$ at 25 degrees Celsius, a pH of 9 gives a pOH of \(14-9=5\). "9" is wrong because that would imply pH and pOH are equal, which only occurs at neutral pH 7. Remembering this constant-sum relationship allows quick conversion between pH and pOH values.

Phenolphthalein transitions from colorless to pink in the pH range of about 8 to 10, which closely matches the slightly basic equivalence point of a strong acid-strong base titration where the salt formed is neutral. "It remains the same color throughout all pH ranges" is wrong because indicators are chosen specifically because their color changes over a defined pH range. Selecting an indicator whose color change range overlaps the equivalence point pH is essential for accurate titration results.

At the equivalence point of a strong acid-weak base titration, the product is the conjugate acid of the weak base, which hydrolyzes water to release extra \(H^+\) ions, making the solution slightly acidic. "Exactly 7, because acid and base always neutralize to neutral" is wrong because neutrality at the equivalence point only occurs when both the acid and base are strong. Predicting equivalence point pH requires considering the strength of the acid and base involved, not just assuming neutrality.

The autoionization of water, \(H_2O \rightleftharpoons H^+ + OH^-\), is an endothermic process, so according to Le Chatelier's principle, raising the temperature shifts the equilibrium toward more product ions and increases \(K_w\). "Because temperature has no effect on equilibrium constants" is wrong because equilibrium constants, including \(K_w\), are inherently temperature-dependent. This means neutral pH is not always exactly 7 at temperatures other than 25 degrees Celsius, even though \([H^+]=[OH^-]\) still holds.

$HCl$ is a strong acid that ionizes completely, while $CH_3COOH$ (acetic acid) is a moderately weak acid, and \(H_2CO_3\) (carbonic acid) is an even weaker acid with a smaller ionization constant, giving the order $HCl$ > $CH_3COOH$ > \(H_2CO_3\). "All three ionize to the same extent" is wrong because these three acids have very different \(K_a\) values and degrees of dissociation in water. Comparing relative acid strength through ionization constants is a key analytical skill for predicting reaction behavior.

The steep vertical jump on a titration curve occurs because near the equivalence point, only trace amounts of unreacted acid or base remain, so even a tiny addition of titrant causes a dramatic shift in \([H^+]\) and thus pH. "The solution becomes a buffer at that pH" is wrong because a buffer requires significant amounts of both a weak acid and its conjugate base, which is not the case at the sharp equivalence point of a strong acid-strong base titration. Recognizing this steep region helps students select an appropriate indicator and interpret titration curves accurately.

Carbonic acid exists in a reversible equilibrium with bicarbonate ion ($HCO_3^-$), and this conjugate acid-base pair can absorb added \(H^+\) or \(OH^-\) to resist large pH swings in blood plasma. "It converts instantly and irreversibly to carbon dioxide gas" is wrong because the carbonic acid-bicarbonate equilibrium is reversible and dynamically responsive, not a one-way irreversible process. This buffer system illustrates how weak acid-conjugate base pairs maintain physiological pH stability.

When a weak acid is diluted, its ionization equilibrium shifts to produce relatively more ions to partially compensate, so its pH changes less than expected from dilution alone, whereas a strong acid's pH shifts more predictably following straightforward dilution math. "Both acids show identical pH changes upon dilution" is wrong because equilibrium shifts uniquely affect weak acids while strong acids remain fully ionized regardless of concentration. Understanding Le Chatelier's principle applied to weak acid equilibria explains this differing dilution behavior.

Since $pH = 14 - pOH = 11$, and \([H^+] = 10^{-pH}\), the hydrogen ion concentration is \(1\times10^{-11}\) M. "\(1\times10^{-3}\) M" is wrong because that value actually corresponds to the hydroxide ion concentration, $[OH^-]=10^{-pOH}$, not \([H^+]\). Students must carefully distinguish between pH-based and pOH-based concentration calculations to avoid mixing up \(H^+\) and \(OH^-\).

For a weak acid where ionization is small, the equilibrium approximation gives \([H^+] \approx \sqrt{K_a \times C}\), derived from setting \(K_a = x^2/C\) and solving for \(x\). "\(K_a \times C\)" is wrong because that expression does not follow from the correct rearrangement of the equilibrium expression and would give a numerically incorrect concentration. This square-root approximation is a standard tool for solving weak acid equilibrium problems when the acid ionizes only slightly.

At the equivalence point of a weak acid-strong base titration, the product is the conjugate base of the weak acid, which reacts with water to generate extra \(OH^-\) ions through hydrolysis, pushing the pH above 7. "The equivalence point always equals the \(pK_a\) of the acid" is wrong because the \(pK_a\) actually corresponds to the half-equivalence point, where \([HA]=[A^-]\), not the full equivalence point. Understanding hydrolysis of the conjugate base is essential for correctly predicting equivalence point pH in these titrations.

The Henderson-Hasselbalch equation states \(pH = pK_a + \log([A^-]/[HA])\), and when \([A^-]=[HA]\), the logarithmic term equals zero, so pH equals \(pK_a\) exactly. "Equal to 7 regardless of the acid used" is wrong because the resulting pH depends entirely on the specific \(pK_a\) of the weak acid chosen, not a fixed neutral value. This equal-concentration condition is a key reference point for designing buffers at a target pH.

A larger \(K_a\) value indicates a stronger tendency to ionize, so \(HA\) releases more \(H^+\) ions at equal concentration than \(HB\), resulting in a lower, more acidic pH. "Both acids produce identical pH values since concentrations are equal" is wrong because pH depends on the actual \(H^+\) concentration produced through ionization, which differs based on acid strength even at equal starting concentrations. Comparing \(K_a\) values, not just initial concentration, is essential for predicting relative acidity among weak acids.

After the first proton is removed, the resulting ion carries a negative charge that more strongly attracts the remaining proton, making it harder to remove and resulting in a smaller \(K_{a2}\) compared to \(K_{a1}\). "The second proton is chemically identical and should ionize equally" is wrong because the electrostatic environment changes significantly after the first ionization, altering the ease of removing the second proton. This general trend of decreasing successive ionization constants applies to all polyprotic acids.

Since \(pH = pK_a + \log([A^-]/[HA])\), and here \([A^-]/[HA] = 0.10/0.20 < 1\), the logarithm is negative, making the pH lower than the \(pK_a\) value of 4.74. "The pH cannot be estimated without knowing total volume" is wrong because the Henderson-Hasselbalch equation depends only on the ratio of moles (or concentrations) of conjugate base to acid, and volume cancels out in that ratio. This ratio-based reasoning allows quick pH estimation for buffers without needing absolute concentrations.

A polyprotic acid like \(H_3PO_4\) has three ionizable protons that are neutralized in a stepwise fashion as base is added, each corresponding to its own distinct equivalence point on the titration curve reflecting a different \(K_a\). "Polyprotic acids release all protons simultaneously at one equivalence point" is wrong because the differing ionization constants for each proton mean they are neutralized at different points, not all at once. Recognizing stepwise neutralization is key to interpreting complex titration curves involving polyprotic species.

Taking \(pH = -\log(4.5\times10^{-6})\) gives approximately \(5.35\), calculated by summing \(-\log(4.5) \approx -0.65\) and \(-\log(10^{-6}) = 6\), yielding \(6 - 0.65 = 5.35\). "4.50" is wrong because it incorrectly treats the exponent alone as the pH without accounting for the coefficient 4.5 in the concentration. Precise pH calculations require handling both the coefficient and the exponent of scientific notation using logarithm rules.

In real laboratory conditions, dissolved atmospheric carbon dioxide can form small amounts of carbonic acid, and measurement uncertainty in glassware or pH meters can cause the observed pH to deviate slightly from the theoretical value of exactly 7. "Strong acids and bases never react completely in reality" is wrong because strong acids and bases are defined by their essentially complete ionization and reaction with each other under normal conditions. Understanding the gap between theoretical stoichiometric predictions and real experimental results is important for interpreting lab data critically.

Adding the conjugate base salt increases the concentration of \(A^-\) already present in the ionization equilibrium, and by Le Chatelier's principle, this shift suppresses further ionization of the weak acid \(HA\), favoring the reverse reaction toward the unionized form. "It converts the weak acid into a strong acid" is wrong because adding a common ion does not alter the fundamental ionization constant \(K_a\) or the intrinsic strength of the acid. The common ion effect is the underlying principle that explains why buffers resist pH change and how conjugate base additions shift weak acid equilibria.

Despite fluorine's high electronegativity, the H-F bond is unusually strong and short, and significant ion-pairing and hydrogen bonding effects in solution prevent complete dissociation, making \(HF\) a weak acid unlike the readily dissociating $HCl\(. "\)HF$ has a smaller molecular size, which always indicates stronger acids" is wrong because molecular size alone does not determine acid strength, and this generalization contradicts the actual weak behavior of \(HF\). This example illustrates that bond strength and solvation effects, not electronegativity alone, determine overall acid strength trends among the hydrohalic acids.

Between the first and second equivalence points, significant amounts of the singly deprotonated acid species and its further deprotonated conjugate base coexist, forming a buffer pair that resists sharp pH changes as more base is added. "The pH increases in a perfectly linear fashion with no resistance to change" is wrong because titration curves show flattened buffering regions rather than a smooth linear increase, reflecting the resistance provided by the conjugate pair. Recognizing these buffering plateaus helps students correctly interpret the shape of polyprotic acid titration curves.

A Brønsted-Lowry base is defined as a species that accepts a proton (\(H^+\)) from another substance, forming a conjugate acid in the process. The choice "A proton donor" describes a Brønsted-Lowry acid, not a base, since acids release \(H^+\) ions to the base. Students should remember that Brønsted-Lowry theory focuses on proton transfer, which broadens the definition of acids and bases beyond Arrhenius theory to include reactions in nonaqueous solvents.

Since $HCl$ and $NaOH$ react in a \(1:1\) mole ratio and both solutions have the same molarity (\(0.100\ M\)), the volume of base needed to neutralize the acid equals the volume of acid, so \(25.0\ mL\) of $NaOH$ supplies the same $0.00250\ mol$ as the acid. The choice "\(12.5\ mL\)" is incorrect because it would only deliver half the moles of $NaOH$ needed to match the $HCl$ present, leaving excess acid unneutralized. In neutralization stoichiometry, students must always balance moles of \(H^+\) and \(OH^-\) using the reaction ratio and concentrations, not assume equal volumes are needed unless the stoichiometric ratio and molarities are equal.