Acids and Bases: Key Concepts | 酸与碱：考点精讲

路易斯理论进一步扩展了定义：酸是电子对受体，碱是电子对供体。这解释了不涉及质子的反应，例如BF₃ + NH₃ → F₃B-NH₃，其中BF₃接受来自NH₃的电子对。

This theory is particularly useful in organic and coordination chemistry, but for most acid-base problems in IB and Edexcel, the Brønsted-Lowry theory is sufficient.

该理论在有机和配位化学中特别有用，但对于IB和Edexcel中的大多数酸碱问题，布朗斯特-劳里理论已足够。

4. Strong and Weak Acids and Bases | 强酸与弱酸、强碱与弱碱

Strong acids completely dissociate in water, such as HCl, H₂SO₄ (first dissociation), and HNO₃. Their dissociation is essentially one-way: HCl → H⁺ + Cl⁻. Weak acids, like CH₃COOH, partially dissociate, establishing an equilibrium: CH₃COOH ⇌ H⁺ + CH₃COO⁻.

强酸在水中完全解离，例如HCl、H₂SO₄（第一步解离）和HNO₃。其解离基本上是单向的：HCl → H⁺ + Cl⁻。弱酸，如CH₃COOH，部分解离，建立平衡：CH₃COOH ⇌ H⁺ + CH₃COO⁻。

Similarly, strong bases like NaOH and KOH fully dissociate to give OH⁻, while weak bases like NH₃ only partially react with water to produce OH⁻.

类似地，强碱如NaOH和KOH完全解离产生OH⁻，而弱碱如NH₃仅部分与水反应生成OH⁻。

5. pH and pOH Calculations | pH与pOH计算

The pH is defined as pH = -log₁₀[H⁺], and pOH = -log₁₀[OH⁻]. In water at 25°C, pH + pOH = 14. For a strong monoprotic acid of concentration c, [H⁺] = c, so pH = -log₁₀c.

pH定义为pH = -log₁₀[H⁺]，pOH = -log₁₀[OH⁻]。在25°C的水中，pH + pOH = 14。对于浓度为c的强一元酸，[H⁺] = c，因此pH = -log₁₀c。

For weak acids, the acid dissociation constant Kₐ is used: Kₐ = [H⁺][A⁻] / [HA]. For a weak acid of initial concentration c and very small dissociation, [H⁺] ≈ √(Kₐc). Then pH = -log₁₀[H⁺].

对于弱酸，使用酸解离常数Kₐ：Kₐ = [H⁺][A⁻] / [HA]。对于初始浓度为c且解离度很小的弱酸，[H⁺] ≈ √(Kₐc)。然后pH = -log₁₀[H⁺]。

6. The Ionic Product of Water, Kₑ | 水的离子积 Kₑ

Water undergoes autoionization: 2H₂O ⇌ H₃O⁺ + OH⁻. The equilibrium constant is Kₑ = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C. This value changes with temperature; at higher temperatures Kₑ increases, making neutral pH different from 7.

水发生自耦解离：2H₂O ⇌ H₃O⁺ + OH⁻。平衡常数Kₑ = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴（25°C）。该值随温度变化；较高温度下Kₑ增大，使得中性pH不等于7。

This concept is essential when calculating the pH of very dilute solutions or strong bases, where the contribution of H⁺ from water cannot be ignored.

这一概念在计算极稀溶液或强碱的pH时至关重要，此时不能忽略来自水的H⁺贡献。

7. Acid-Base Titrations and pH Curves | 酸碱滴定与pH曲线

Titration curves plot pH against the volume of titrant added. For a strong acid-strong base titration, the equivalence point is at pH 7. For a weak acid-strong base titration, the equivalence point is above pH 7 due to the formation of the conjugate base.

滴定曲线描绘了pH随滴定剂加入体积的变化。强酸-强碱滴定的等当点在pH 7。弱酸-强碱滴定的等当点高于7，由于生成了共轭碱。

The buffer region occurs where the pH changes slowly, typically around the half-equivalence point for weak acid or weak base titrations. Indicators with appropriate pKₐ values are chosen so their colour change coincides with the steep part of the curve.

缓冲区域出现在pH变化缓慢的地方，通常在弱酸或弱碱滴定的半等当点附近。选择具有适当pKₐ值的指示剂，使其颜色变化与曲线的陡峭部分重合。

8. Buffer Solutions | 缓冲溶液

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 (e.g., CH₃COOH/CH₃COO⁻) or a weak base and its conjugate acid (e.g., NH₃/NH₄⁺).

缓冲溶液能抵抗因少量酸或碱加入而引起的pH变化。它由弱酸及其共轭碱（例如CH₃COOH/CH₃COO⁻）或弱碱及其共轭酸（例如NH₃/NH₄⁺）组成。

The Henderson-Hasselbalch equation is used to calculate the pH of a buffer: pH = pKₐ + log₁₀([A⁻]/[HA]), where [A⁻] is the concentration of conjugate base and [HA] is the concentration of the weak acid.

亨德森-哈塞尔巴尔赫方程用于计算缓冲液的pH：pH = pKₐ + log₁₀([A⁻]/[HA])，其中[A⁻]是共轭碱浓度，[HA]是弱酸浓度。

Buffers are crucial in biological systems, such as maintaining blood pH around 7.4 using the carbonic acid/bicarbonate system.

缓冲液在生物系统中至关重要，例如利用碳酸/碳酸氢盐系统维持血液pH约7.4。

9. Acid Deposition and Environmental Context | 酸沉降与环境背景

Acid deposition, including acid rain, is caused by SO₂ and NOₓ emissions reacting with water in the atmosphere to form H₂SO₄ and HNO₃. This links to acid-base chemistry through the formation of acidic solutions and their impact on ecosystems.

酸沉降，包括酸雨，是由SO₂和NOₓ排放与大气中的水反应生成H₂SO₄和HNO₃引起的。这通过酸性溶液的形成及其对生态系统的影响与酸碱化学联系起来。

Exam questions may ask about the equations involved: for example, 2SO₂ + O₂ + 2H₂O → 2H₂SO₄, or the neutralization of acidified lakes using limestone (CaCO₃).

考试可能涉及相关方程式：例如，2SO₂ + O₂ + 2H₂O → 2H₂SO₄，或使用石灰石（CaCO₃）中和酸化湖泊。

10. Common Exam Pitfalls and Tips | 常见考试陷阱与提示

Students often confuse strength and concentration; a strong acid is fully dissociated regardless of its concentration, while a weak acid is only partially dissociated even at high concentration. Be careful when calculating pH for diprotic acids like H₂SO₄, as the second dissociation may be partial.

学生常混淆强度与浓度；强酸无论浓度如何都完全解离，而弱酸即使高浓度也仅部分解离。计算二元酸如H₂SO₄的pH时要小心，因为第二步解离可能不完全。

Always check the temperature when using Kₑ = 1.0 × 10⁻¹⁴, and remember that pH + pOH = 14 only at 25°C. In buffer calculations, ensure you use equilibrium concentrations, not initial amounts.

使用Kₑ = 1.0 × 10⁻¹⁴时务必核对温度，记住pH + pOH = 14仅在25°C成立。在缓冲计算中，确保使用平衡浓度而非初始量。

For indicator selection, the pKₐ of the indicator should be within ±1 of the equivalence point pH, and the colour change should be distinct.

指示剂选择方面，指示剂的pKₐ应在等当点pH的±1范围内，且颜色变化应明显。

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