Comprehensive Guide to Acids, Bases, Salts, and pH

Key Terms and Principles of Acid Strength

• Acid Strength Determination

  • Acid strength is governed by the degree of ionization or dissociation of the acid molecules when dissolved in water, rather than the total concentration of the solution.
  • A strong acid dissociates completely ($100\%$) into its constituent ions in an aqueous solution. For example, hydrochloric acid completely ionizes: $HCl_{(aq)} \rightarrow H^+_{(aq)} + Cl^-_{(aq)}$.
  • A weak acid dissociates only partially in water, establishing a dynamic chemical equilibrium between the intact molecule and the dissociated ions. For example, acetic acid partially ionizes: $CH_3COOH_{(aq)} \rightleftharpoons H^+_{(aq)} + CH_3COO^-_{(aq)}$.
  • Strong acids possess weak conjugate bases with low affinity for protons, facilitating complete dissociation. Weak acids possess relatively stronger conjugate bases that readily recombine with hydrogen ions.

• Foundational Vocabulary and Definitions

  • Acid: A chemical substance that acts as a proton ($H^+$) donor in aqueous solutions, generating hydronium ions ($H_3O^+$) and possessing a $\text{pH}$ value less than $7.0$.
  • Base: A chemical substance that acts as a proton ($H^+$) acceptor or yields hydroxide ions ($OH^-$) in aqueous solution, exhibiting a $\text{pH}$ value greater than $7.0$.
  • Ion: An atom or molecule that has acquired a net electrical charge due to the loss or gain of one or more valence electrons (yielding positively charged cations or negatively charged anions).
  • Hydronium Ion ($H_3O^+$): The positive ion formed when a released hydrogen ion ($H^+$) covalently bonds with a water molecule ($H_2O$).
  • Hydroxide Ion ($OH^-$): A diatomic anion consisting of oxygen and hydrogen with a negative charge ($OH^-$), characteristic of basic solutions.
  • Salt: An ionic chemical compound composed of a cation (derived from a base) and an anion (derived from an acid), formed typically through a neutralization reaction.
  • Neutralization Reaction: A double-replacement chemical reaction between an acid and a base producing a salt and water: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + H_2O$.
  • $\text{pH}$: A logarithmic scale measuring the concentration of hydrogen ions ($[H^+]$) in an aqueous solution: $\text{pH} = -\log_{10}([H^+])$.
  • Electrolyte: A compound that dissociates into free ions when dissolved in water or melted, producing an electrically conductive medium.

Everyday Acids and Bases

• Common Household and Biological Acids

  • Citric Acid ($C_6H_8O_7$): Found naturally in citrus fruits, including lemons, limes, and oranges; utilized as a flavor enhancer and preservative.
  • Acetic Acid ($CH_3COOH$): The primary acidic constituent of household vinegar (typically $4\%$ to $8\%$ acetic acid by volume), used in food preparation and cleaning.
  • Hydrochloric Acid ($HCl$): Present in gastric juice within the stomach at concentrations sufficient to maintain a gastric $\text{pH}$ of $1.5$ to $3.5$ for protein digestion and pathogen destruction; commercially sold as muriatic acid.
  • Carbonic Acid ($H_2CO_3$): Formed when carbon dioxide dissolves in water; present in all carbonated soft drinks and beverages.
  • Lactic Acid ($C_3H_6O_3$): Present in sour milk, yogurt, and generated in animal muscle tissues during anaerobic respiration.

• Common Household and Biological Bases

  • Sodium Hydroxide ($NaOH$): Also known as lye or caustic soda; a potent strong base utilized in commercial drain openers, oven cleaners, and soap manufacturing.
  • Sodium Bicarbonate ($NaHCO_3$): Commonly known as baking soda; a mild base used as a leavening agent in baking and as an antacid to relieve heartburn.
  • Magnesium Hydroxide ($Mg(OH)_2$): Formulated as milk of magnesia; used medically as an antacid to neutralize excess stomach acidity and as an osmotic laxative.
  • Ammonia ($NH_3$): A weak base commonly formulated into aqueous household glass and surface cleaners.
  • Calcium Hydroxide ($Ca(OH)_2$): Known as slaked lime; utilized in water treatment, mortar formulation, and agriculture to reduce soil acidity.

• Behavior and Role of Ions in Acid-Base Chemistry

  • Pure liquid water undergoes self-ionization (autoionization) to a minute degree: $2H_2O_{(l)} \rightleftharpoons H_3O^+_{(aq)} + OH^-_{(aq)}$.
  • In neutral water at $25\,^{\circ}\text{C}$, the concentrations of hydronium and hydroxide ions are precisely equal: $[H_3O^+] = [OH^-] = 1.0 \times 10^{-7}\,mol\,dm^{-3}$.
  • When an acid dissolves, it contributes additional $H^+$ (or $H_3O^+$) ions to the solution, shifting the equilibrium such that $[H^+] > [OH^-]$.
  • When a base dissolves, it either directly dissociates to release $OH^-$ ions or abstracts protons from water molecules to generate $OH^-$ ions, shifting the balance such that $[OH^-] > [H^+]$.

• Physiological Regulation of Blood $\text{pH}$

  • Normal arterial human blood $\text{pH}$ must remain within an extraordinarily narrow physiological range of $7.35$ to $7.45$, with an optimum average of $7.40$.
  • Severe physiological distress, metabolic dysfunction, or death occurs if systemic blood $\text{pH}$ drops below $7.35$ (acidosis) or rises above $7.45$ (alkalosis).
  • The primary buffer system regulating blood $\text{pH}$ is the carbonic acid-bicarbonate buffer pair:     $CO_{2(aq)} + H_2O_{(l)} \rightleftharpoons H_2CO_{3(aq)} \rightleftharpoons H^+_{(aq)} + HCO_{3(aq)}^-$
  • Respiratory Regulation: Changes in ventilation rate alter blood carbon dioxide levels. When excess $H^+$ accumulates, respiration increases to exhale $CO_2$, driving the equilibrium to the left and raising blood $\text{pH}$. Hypoventilation retains $CO_2$, driving equilibrium rightward and lowering blood $\text{pH}$.
  • Renal Regulation: The kidneys provide long-term regulation by selectively excreting excess hydrogen ions ($H^+$) into the urine while reabsorbing or synthesizing bicarbonate ions ($HCO_3^-$) back into the bloodstream.

• Sources and Environmental Impacts of Acid Rain

  • Sources of Acidic Precursors:
  • Combustion of sulfur-bearing fossil fuels (primarily coal in power generation) emits sulfur dioxide ($SO_2$).
  • High-temperature combustion in motor vehicle internal combustion engines and industrial boilers oxidizes atmospheric nitrogen to form nitrogen monoxide ($NO$) and nitrogen dioxide ($NO_2$), collectively termed $NO_x$.
  • Atmospheric Chemical Transformations:
  • Sulfur dioxide oxidizes and combines with atmospheric moisture to form sulfurous and sulfuric acid: $SO_2 + H_2O \rightarrow H_2SO_3$ and $2SO_2 + O_2 + 2H_2O \rightarrow 2H_2SO_4$.
  • Nitrogen dioxide reacts with water vapor to form nitric and nitrous acids: $2NO_2 + H_2O \rightarrow HNO_3 + HNO_2$.
  • Environmental Impacts:
  • Aquatic Ecosystems: Unpolluted rain has a natural $\text{pH}$ of approximately $5.6$ due to dissolved atmospheric $CO_2$; acid precipitation often reaches $\text{pH}$ values of $4.0$ to $4.5$ or lower. This acidifies lakes, ponds, and streams, killing aquatic invertebrates and fish eggs.
  • Soil Chemistry and Flora: Acid rain mobilizes and leaches toxic aluminum ions ($Al^{3+}$) from soil minerals into runoff, while depleting vital plant nutrients such as calcium ($Ca^{2+}$) and magnesium ($Mg^{2+}$), stunting forest growth and damaging protective leaf cuticles.
  • Built Infrastructure: Corrodes architectural structures, metal bridges, and historical monuments composed of limestone or marble by dissolving calcium carbonate: $CaCO_{3(s)} + H_2SO_{4(aq)} \rightarrow CaSO_{4(aq)} + H_2O_{(l)} + CO_{2(g)}$.

Characteristics and Properties of Acids, Bases, and Salts

• Comparative Properties of Acids and Bases

  • Taste:
  • Acids exhibit a distinct sharp, sour taste (characteristic of citric acid in citrus fruits and acetic acid in vinegar).
  • Bases possess an astringent, bitter taste (characteristic of tonic water containing quinine, unsweetened cocoa, or soap).
  • Tactile Texture (Feel):
  • Acids feel watery between the fingers; concentrated acids cause severe burning and tissue damage.
  • Bases feel distinctly slippery or soapy to the touch because they react with fatty acids and natural oils on human skin in a process called saponification.
  • Dominant Ion Concentrations:
  • Acids contain an elevated concentration of hydrogen/hydronium ions ($[H^+] > [OH^-]$), with $[H^+] > 1.0 \times 10^{-7}\,mol\,dm^{-3}$ at $25\,^{\circ}\text{C}$.
  • Bases contain an elevated concentration of hydroxide ions ($[OH^-] > [H^+]$), with $[OH^-] > 1.0 \times 10^{-7}\,mol\,dm^{-3}$ at $25\,^{\circ}\text{C}$.
  • Chemical Reactivity:
  • Reactivity of Acids:
    • React vigorously with active metals located above hydrogen in the activity series (such as magnesium, zinc, and iron) via single-replacement reactions to produce hydrogen gas and a salt:         $Mg_{(s)} + 2HCl_{(aq)} \rightarrow MgCl_{2(aq)} + H_{2(g)}$
    • React with metal carbonates and bicarbonates to generate carbon dioxide gas, water, and an ionic salt:         $CaCO_{3(s)} + 2HCl_{(aq)} \rightarrow CaCl_{2(aq)} + H_2O_{(l)} + CO_{2(g)}$
    • React stoichiometrically with bases in neutralization reactions to produce a salt and liquid water.
    • Turn blue litmus paper red.
  • Reactivity of Bases:
    • Do not react with most structural metals under standard conditions (with exceptions such as amphoteric metals like aluminum or zinc under highly concentrated basic conditions).
    • React with acids in neutralization reactions to neutralize acidity.
    • React with fats, lipids, and organic matter (useful in saponification reactions and chemical drain unblocking).
    • Turn red litmus paper blue.

• Nature and Characteristics of Salts

  • Definition: A salt is an ionic compound composed of an array of cations and anions held together by strong electrostatic forces of attraction (ionic bonds), produced by the reaction of an acid with a base.
  • Physical Properties of Salts:
  • Exist as crystalline solids at room temperature ($20\,^{\circ}\text{C}$ to $25\,^{\circ}\text{C}$) with orderly three-dimensional crystal lattice structures.
  • Possess high melting points and high boiling points due to the extensive thermal energy required to overcome strong lattice energies (e.g., sodium chloride melts at $801\,^{\circ}\text{C}$).
  • Display hardness and brittleness; mechanical stress aligns ions of identical charge adjacent to each other, producing repulsive forces that cause crystal cleavage.
  • Function as electrical insulators in the solid state because ions are fixed rigidly in the crystal lattice.
  • Function as strong electrical conductors (electrolytes) when dissolved in water or melted into a liquid state, as the ions dissociate and move freely.
  • Chemical Properties of Salts:
  • Can be neutral, acidic, or basic in aqueous solution depending on the relative strengths of the parent acid and parent base:
    • Neutral Salt: Formed from a strong acid and a strong base (e.g., $HCl + NaOH \rightarrow NaCl + H_2O$; solution $\text{pH} = 7.0$).
    • Acidic Salt: Formed from a strong acid and a weak base (e.g., $HCl + NH_3 \rightarrow NH_4Cl$; the ammonium ion undergoes hydrolysis to yield an acidic solution with $\text{pH} < 7.0$).
    • Basic Salt: Formed from a weak acid and a strong base (e.g., $CH_3COOH + NaOH \rightarrow CH_3COONa + H_2O$; the acetate ion undergoes hydrolysis to yield a basic solution with $\text{pH} > 7.0$).
  • Participate in double-replacement (precipitation) reactions when paired with compatible aqueous ionic reagents to form insoluble precipitates.

Understanding the $\text{pH}$ Scale

• Principles and Mathematics of $\text{pH}$ Measurement

  • The term $\text{pH}$ stands for the "potential of hydrogen" or "power of hydrogen."
  • The scale is inverse and logarithmic: each single-unit change on the $\text{pH}$ scale represents a tenfold ($10\times$) difference in hydrogen ion concentration ($[H^+]$).
  • Mathematical formula for $\text{pH}$:     $\text{pH} = -\log_{10}([H^+])$
  • Mathematical formula for $\text{pOH}$:     $\text{pOH} = -\log_{10}([OH^-])$
  • Autoionization equilibrium constant for water at $25\,^{\circ}\text{C}$ ($K_w$):     $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$
  • The relationship linking $\text{pH}$ and $\text{pOH}$ at $25\,^{\circ}\text{C}$ is constant:     $\text{pH} + \text{pOH} = 14$

• $\text{pH}$ Scale Numerical Classifications

  • Acidic Solutions:
  • Range: $\text{pH} < 7.0$.
  • Strong acids reside typically in the range of $\text{pH} = 0.0$ to $3.0$ (e.g., $1.0\,mol\,dm^{-3}\,HCl$ has a $\text{pH} = 0.0$, stomach acid $\text{pH} \approx 1.5$, lemon juice $\text{pH} \approx 2.0$).
  • Weak acids reside typically in the range of $\text{pH} = 3.0$ to $6.9$ (e.g., coffee $\text{pH} \approx 5.0$, unpolluted rain $\text{pH} \approx 5.6$, milk $\text{pH} \approx 6.5$).
  • Neutral Solutions:
  • Defined as $\text{pH} = 7.0$ at $25\,^{\circ}\text{C}$.
  • Hydrogen and hydroxide ion concentrations are exactly equal: $[H^+] = [OH^-] = 1.0 \times 10^{-7}\,mol\,dm^{-3}$ (e.g., pure deionized water).
  • Basic (Alkaline) Solutions:
  • Range: $\text{pH} > 7.0$.
  • Weak bases reside typically in the range of $\text{pH} = 7.1$ to $10.5$ (e.g., human blood $\text{pH} \approx 7.4$, sea water $\text{pH} \approx 8.1$, baking soda solution $\text{pH} \approx 8.5$, milk of magnesia $\text{pH} \approx 10.5$).
  • Strong bases reside typically in the range of $\text{pH} = 11.0$ to $14.0$ (e.g., household ammonia $\text{pH} \approx 11.5$, liquid bleach $\text{pH} \approx 12.5$, $1.0\,mol\,dm^{-3}\,NaOH$ has a $\text{pH} = 14.0$).

• Dynamic Shifts in Ion Concentrations and $\text{pH}$

  • Effects of Changing $[H^+]$:
  • As $[H^+]$ increases, the $\text{pH}$ decreases proportionately (the solution becomes more acidic).
  • As $[H^+]$ decreases, the $\text{pH}$ increases proportionately (the solution becomes less acidic and more basic).
  • A decrease of $1$ unit of $\text{pH}$ represents a $10\text{-fold}$ increase in $[H^+]$; a decrease of $2$ units represents a $100\text{-fold}$ increase in $[H^+]$.
  • Effects of Changing $[OH^-]$:
  • As $[OH^-]$ increases, $[H^+]$ must decrease to satisfy $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$; consequently, $\text{pH}$ increases (the solution becomes more basic/alkaline).
  • As $[OH^-]$ decreases, $[H^+]$ increases correspondingly, causing the $\text{pH}$ to decrease (the solution becomes less basic and more acidic).

Scientific Review and Synthesis

• Structured Review Process

  • Review all recorded properties, reactions, and formulas for two uninterrupted minutes.
  • Focus specifically on the relationships between ion dominance ($H^+$ versus $OH^-$), conjugate pairs, and logarithmic $\text{pH}$ shifts.

• Concise Synthesis Summary

  • Acids yield $H^+$ and lower $\text{pH}$, bases increase $OH^-$, and neutralization forms ionic salts and water.