📚 A-Level Chemistry: Mastering pH Calculations | A-Level 化学:pH计算 考点精讲
pH calculations are a fundamental part of A-Level Chemistry, bridging equilibrium theory, acid-base behaviour, and practical titration analysis. Mastering these calculations involves understanding the ionic product of water, strong and weak acids and bases, buffer systems, and the interpretation of titration curves. This article provides a comprehensive revision guide to all key concepts and problem types required for A-Level examinations.
pH计算是A-Level化学的基础部分,连接了平衡理论、酸碱行为以及滴定分析实操。掌握这些计算需要理解水的离子积、强弱酸和强弱碱、缓冲体系以及滴定曲线的解读。本文针对A-Level考试所需的所有关键概念和题型,提供一份全面的复习指南。
1. Introduction to pH and the pH Scale | pH和pH标度简介
The pH scale is a logarithmic measure of hydrogen ion concentration in aqueous solution at a given temperature. It is defined as:
pH标度是在特定温度下,水溶液中氢离子浓度的对数度量。其定义为:
pH = –log[H⁺]
where [H⁺] is the concentration of hydrogen ions in mol dm⁻³, and the logarithm is base 10. A solution with pH 7 at 25°C is considered neutral, because [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ mol dm⁻³. As [H⁺] increases, pH decreases; a change of one pH unit corresponds to a ten-fold change in hydrogen ion concentration. Most laboratory measurements use pH meters, but calculations rely on the relationship above and the ionic product of water.
其中[H⁺]是氢离子浓度,单位为mol dm⁻³,对数的底为10。25°C时,pH=7的溶液被视为中性,因为[H⁺] = [OH⁻] = 1.0 × 10⁻⁷ mol dm⁻³。随着[H⁺]增大,pH减小;pH值每改变1个单位,氢离子浓度发生10倍的变化。大多数实验室测量使用pH计,但计算依赖于上述关系及水的离子积。
2. The Ionic Product of Water, Kw | 水的离子积Kw
Water undergoes self-ionisation to a very small extent, establishing an equilibrium that is crucial for all aqueous acid-base calculations:
水会发生极微弱的自耦电离,建立对一切水溶液酸碱计算都至关重要的平衡:
2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)
The ionic product of water, Kw, is defined as:
水的离子积Kw定义为:
Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol² dm⁻⁶ (at 298 K)
Note that [H⁺] is often used interchangeably with [H₃O⁺]. Kw has a fixed value at a given temperature; it increases with rising temperature because the forward reaction is endothermic. This expression allows you to calculate [OH⁻] from a known [H⁺] or vice versa, which is essential when working with strong bases and neutralisation reactions.
注意[H⁺]通常可互换地表示[H₃O⁺]。Kw在给定温度下为定值;因为正反应吸热,其值随温度升高而增大。利用该表达式,可由已知的[H⁺]求[OH⁻],反之亦然,这对于处理强碱和中和反应至关重要。
3. Calculating pH of Strong Acids | 强酸的pH计算
Strong acids such as HCl, HNO₃, and H₂SO₄ (first dissociation) are assumed to dissociate completely in dilute aqueous solution. Therefore, the concentration of H⁺ is equal to the initial concentration of the acid, after accounting for the stoichiometry.
强酸如HCl、HNO₃以及H₂SO₄(一级解离)在稀水溶液中可视为完全解离。因此,考虑到化学计量比后,H⁺浓度等于酸的初始浓度。
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For a monoprotic strong acid HA: [H⁺] = c(acid).
对于一元强酸HA:[H⁺] = c(酸)。
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For diprotic strong acid H₂SO₄, the first dissociation is complete, giving [H⁺] = c(acid). The second dissociation is partial but is often treated as complete at A-Level for the purpose of calculating total [H⁺] from the first proton, but careful exam questions may ask for the contribution from the second proton (which requires knowledge of Kₐ₂). Usually, at this level, you simply take [H⁺] from the fully dissociated first proton as the main contributor, unless stated otherwise.
对于二元强酸H₂SO₄,一级解离完全,[H⁺] = c(酸)。二级解离部分进行,但在A-Level计算中通常将第一级质子的贡献作为主要H⁺来源,除非题目另有说明。
Once [H⁺] is obtained, pH = –log[H⁺]. Pay attention to significant figures: the number of decimal places in pH should equal the number of significant figures in the concentration.
一旦得到[H⁺],pH = –log[H⁺]。注意有效数字:pH值的小数位数应等于浓度数值的有效数字位数。
4. Calculating pH of Strong Bases | 强碱的pH计算
Strong bases such as NaOH, KOH dissociate completely to release hydroxide ions. For a monoacidic base: [OH⁻] = c(base). To find pH, first calculate pOH and then use the relationship pH + pOH = pKw = 14.00 at 298 K.
强碱如NaOH、KOH完全解离释放氢氧根离子。对于一元碱:[OH⁻] = c(碱)。为求pH,先计算pOH,再使用关系式pH + pOH = pKw = 14.00(298K条件下)。
pOH = –log[OH⁻]
pH = 14 – pOH
For bases such as Ba(OH)₂, which provides two OH⁻ per formula unit, [OH⁻] = 2 × c(salt), provided the base is strong and fully soluble. Always check the dissociation stoichiometry before substituting values.
对于Ba(OH)₂这类每个组成单元提供两个OH⁻的碱,若该碱为强碱且完全溶解,则[OH⁻] = 2 × c(盐)。在代入数值前,务必检查解离的化学计量比。
5. Weak Acids and the Acid Dissociation Constant, Kₐ | 弱酸和酸解离常数Kₐ
Weak acids partially dissociate in water, establishing an equilibrium that is described by the acid dissociation constant Kₐ. For a generic weak acid HA:
弱酸在水中部分解离,建立起由酸解离常数Kₐ描述的平衡。对于通式HA的弱酸:
HA(aq) + H₂O(l) ⇌ H₃O⁺(aq) + A⁻(aq)
Kₐ = [H⁺][A⁻] / [HA]
The magnitude of Kₐ indicates acid strength. The smaller the Kₐ, the weaker the acid. pKₐ = –logKₐ, and a larger pKₐ corresponds to a weaker acid. When carrying out weak acid calculations, two common approximations are used:
Kₐ的大小指示酸的强弱。Kₐ越小,酸越弱。pKₐ = –logKₐ,pKₐ越大对应酸越弱。进行弱酸计算时常使用两个近似假设:
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The concentration of H⁺ from the autoionisation of water is negligible compared to that from the acid.
相较于酸解离产生的H⁺,水自耦电离产生的H⁺可忽略不计。
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The equilibrium concentration of the undissociated acid [HA] is approximately equal to the initial concentration c, because dissociation is small (less than 5% typically).
未解离酸[HA]的平衡浓度近似等于初始浓度c,因为解离度很小(通常小于5%)。
6. Calculating pH of Weak Acids | 弱酸的pH计算
Using the approximation [HA] ≈ c and [H⁺] = [A⁻], the Kₐ expression simplifies to:
利用[HA] ≈ c和[H⁺] = [A⁻],Kₐ表达式简化为:
Kₐ = [H⁺]² / c
Thus:
由此可得:
[H⁺] = √(Kₐ × c)
Then pH = –log[H⁺]. This formula works well when the degree of dissociation is less than about 5%. If the weak acid is not too dilute and Kₐ is small, the approximation is valid. For stronger weak acids or very dilute solutions, a quadratic equation must be solved exactly. Exam questions often ask you to check the validity of the approximation by calculating percent dissociation = ([H⁺]/c) × 100%.
然后 pH = –log[H⁺]。当解离度小于约5%时,该公式效果良好。如果弱酸不太稀且Kₐ较小,近似有效。对于较强的弱酸或极稀溶液,则须严格求解二次方程。考试中常要求通过计算解离百分率 = ([H⁺]/c) × 100% 来检验近似的有效性。
For polyprotic weak acids (e.g., H₂CO₃, H₃PO₄), only the first dissociation contributes significantly to [H⁺] because successive Kₐ values are much smaller. Therefore, treat them as monoprotic weak acids using Kₐ₁.
对于多元弱酸(如H₂CO₃、H₃PO₄),因后续Kₐ值远小于Kₐ₁,仅一级解离对[H⁺]有显著贡献。因此,可将其当作一元弱酸处理,使用Kₐ₁。
7. Weak Bases and the Base Dissociation Constant, Kb | 弱碱和碱解离常数Kb
Weak bases such as ammonia (NH₃) and amines react partially with water to produce hydroxide ions. The equilibrium is described by the base dissociation constant Kb:
弱碱如氨(NH₃)和胺类部分与水反应生成氢氧根离子。该平衡由碱解离常数Kb描述:
NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)
Kb = [NH₄⁺][OH⁻] / [NH₃]
Analogous to weak acids, the smaller the Kb, the weaker the base. pKb = –logKb. For a conjugate acid-base pair, the relationship holds: Kₐ × Kb = Kw, and pKₐ + pKb = 14 at 298 K. This allows conversion between acid and base constants of conjugate species.
类似弱酸,Kb越小,碱越弱。pKb = –logKb。对于共轭酸碱对,存在关系:Kₐ × Kb = Kw,且298K时pKₐ + pKb = 14。这可用于共轭物种酸碱常数之间的换算。
8. Calculating pH of Weak Bases | 弱碱的pH计算
For a weak base with initial concentration c, using the approximation that dissociation is small, we have [OH⁻] = [BH⁺] and [B] ≈ c. Then:
对于初始浓度为c的弱碱,假设解离度很小,则有[OH⁻] = [BH⁺]且[B] ≈ c。于是:
Kb = [OH⁻]² / c
[OH⁻] = √(Kb × c)
Then pOH = –log[OH⁻], and pH = 14 – pOH at 298 K. Just as with weak acids, the approximation must be checked by percent dissociation. Alternatively, if Kₐ of the conjugate acid is given, you can convert using Kb = Kw/Kₐ and solve as above.
然后 pOH = –log[OH⁻],pH = 14 – pOH(298K)。同弱酸一样,需用解离百分率检验近似。若已知共轭酸的Kₐ,也可通过Kb = Kw/Kₐ换算后按上述方法求解。
Weak bases derived from the conjugate base of a weak acid (e.g., CH₃COO⁻) are sometimes encountered in salt hydrolysis, where the anion reacts with water to produce OH⁻, and the pH calculation follows the same weak base approach.
有时会遇到源于弱酸共轭碱的弱碱(例如CH₃COO⁻),即盐水解的情况,阴离子与水反应生成OH⁻,其pH计算遵循同样的弱碱方法。
9. Buffer Solutions and the Henderson-Hasselbalch Equation | 缓冲溶液和Henderson-Hasselbalch方程
A buffer solution resists changes in pH upon addition of small amounts of acid or base. It consists of a weak acid and its conjugate base (acidic buffer) or a weak base and its conjugate acid (basic buffer). The pH of an acidic buffer is given by the Henderson-Hasselbalch equation:
缓冲溶液能抵抗少量外加酸碱带来的pH变化。它由弱酸及其共轭碱(酸性缓冲液)或弱碱及其共轭酸(碱性缓冲液)组成。酸性缓冲液的pH由Henderson-Hasselbalch方程给出:
pH = pKₐ + log([A⁻] / [HA])
This equation is derived from the Kₐ expression and is valid when the concentrations of the acid and its salt are large compared to [H⁺]. It shows that the buffer pH depends primarily on the pKₐ of the weak acid and the ratio of the concentrations of conjugate base to acid.
该方程由Kₐ表达式推导而来,当酸及其盐的浓度远大于[H⁺]时成立。这表明缓冲液的pH主要取决于弱酸的pKₐ以及共轭碱与酸浓度的比值。
When [A⁻] = [HA], pH = pKₐ. Buffering capacity is maximum near this point. You can prepare a buffer by mixing a weak acid with its salt (e.g., CH₃COOH and CH₃COONa) or by partial neutralisation of a weak acid with strong base. Calculations require careful stoichiometry to determine the amounts of acid and conjugate base after mixing.
当[A⁻] = [HA]时,pH = pKₐ,在该点附近缓冲容量最大。可通过将弱酸与其盐混合(如CH₃COOH与CH₃COONa)或通过强碱部分中和弱酸来制备缓冲液。计算时需仔细运用化学计量法确定混合后酸及其共轭碱的量。
10. pH Changes During Titrations – Strong Acid–Strong Base | 滴定过程中的pH变化 – 强酸强碱
A titration curve plots pH against the volume of titre added. For a strong acid–strong base titration, the curve has a characteristic steep vertical region around the equivalence point (pH = 7). Before the equivalence point, pH is determined by the excess strong acid; after, by the excess strong base.
滴定曲线将pH对待加入滴定剂体积作图。强酸-强碱滴定的曲线在等当点(pH=7)附近有一特征性的陡直区域。等当点前,pH由过量的强酸决定;等当点后,由过量的强碱决定。
Calculation steps at any point: determine the moles of H⁺ and OH⁻ initially. Subtract the moles of the limiting reagent. The remaining moles of the excess ion dictate the concentration, considering the total volume. Then calculate [H⁺] or [OH⁻] and convert to pH.
任一点的计算步骤为:确定初始H⁺和OH⁻的物质的量,减去限量试剂的物质的量。过量离子的剩余物质的量除以总体积得到浓度,进而计算[H⁺]或[OH⁻]并转换为pH。
At the equivalence point, the solution contains only the salt (e.g., NaCl) and water, so pH = 7. The steep rise near the equivalence point is due to the logarithmic nature of the pH scale: a tiny excess of base causes a large jump in pH.
在等当点,溶液中仅含盐(如NaCl)和水,因此pH=7。等当点附近的急剧上升源于pH标度的对数特性:极微量的过量碱便可引起pH大幅跃迁。
11. pH Changes During Titrations – Weak Acid–Strong Base | 弱酸强碱滴定
When a weak acid is titrated with a strong base, the curve differs significantly. The initial pH is higher (because the acid is weak), and there is a buffer region where pH changes slowly. The equivalence point pH is greater than 7 due to the hydrolysis of the conjugate base (A⁻) that produces OH⁻.
用强碱滴定弱酸时,曲线明显不同。初始pH较高(因酸为弱酸),并存在一段pH变化缓慢的缓冲区。由于共轭碱(A⁻)水解生成OH⁻,等当点pH大于7。
Before any base is added, the pH is that of the weak acid. In the buffer region (before equivalence), the Henderson-Hasselbalch equation applies. At the half-equivalence point, [HA] = [A⁻], so pH = pKₐ. At the equivalence point, the dominant species is the conjugate base A⁻; calculate its concentration and then treat the solution as a weak base using Kb = Kw/Kₐ.
加入任何碱之前,pH为弱酸的pH。在缓冲区(等当点前),适用Henderson-Hasselbalch方程。在半等当点处,[HA] = [A⁻],故pH = pKₐ。在等当点,主导物种是共轭碱A⁻;计算其浓度,然后将其视为弱碱处理,利用Kb = Kw/Kₐ求解。
After the equivalence point, excess strong base controls the pH, and the calculation is similar to the strong acid–strong base case. The vertical rise is less steep than for strong–strong titrations, and the choice of indicator must account for the higher equivalence pH.
等当点之后,过量的强碱控制pH,计算与强酸强碱情况类似。该曲线垂直上升段不如强强滴定陡峭,且选择指示剂时必须考虑较高的等当点pH。
12. Choosing Indicators for Titrations | 选择滴定指示剂
Acid-base indicators are weak acids or bases whose undissociated and dissociated forms have different colours. Their behaviour is described by the indicator constant KIn and the relationship pH = pKIn + log([In⁻]/[HIn]). The colour change occurs over a range of about pH = pKIn ± 1.
酸碱指示剂本身是弱酸或弱碱,其未解离形与解离形具有不同的颜色。其行为由指示剂常数KIn和关系式pH = pKIn + log([In⁻]/[HIn])描述。颜色变化发生的范围大约在pH = pKIn ± 1。
To choose a suitable indicator for a titration, its pH range must overlap with the steep portion of the titration curve. For strong acid–strong base titrations, indicators with pKIn around 3–10 can be used (e.g., methyl orange, phenolphthalein), because the vertical section spans a wide pH range. For weak acid–strong base titrations, the equivalence point lies in the alkaline region, so an indicator like phenolphthalein (pKIn ≈ 9.3) is suitable. Methyl orange (pKIn ≈ 3.7) changes colour too early and would give a large titration error.
为滴定选择合适的指示剂,其pH变色范围必须与滴定曲线的陡直段重叠。强酸强碱滴定中,pKIn在3–10左右的指示剂均可使用(如甲基橙、酚酞),因为垂直段跨越较宽的pH区间。弱酸强碱滴定的等当点位于碱性区域,因此适合选用酚酞(pKIn ≈ 9.3)这类指示剂。甲基橙(pKIn ≈ 3.7)变色过早,会产生较大的滴定误差。
In weak base–strong acid titrations, the equivalence point is acidic, and methyl orange is appropriate. The exam may ask you to justify indicator choice using the sketch of the titration curve and the indicator’s pKIn.
在弱碱强酸滴定中,等当点为酸性,甲基橙合适。考试可能要求你根据滴定曲线简图和指示剂的pKIn来说明选择理由。
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