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Buffers, Acids/Bases, and Organic vs Inorganic Chemistry — Vocabulary Flashcards

pH, Hydronium, Hydroxide, and the Water Equilibrium

  • pH is a measure of acidity/basicity based on the hydronium ion concentration in solution.

  • Hydronium ion is commonly written as
    [ ext{H}_3 ext{O}^+]

  • The basic relationships:

    • ext{pH} = -\,\log{10} [\text{H}^+] (or [\text{H}3\text{O}^+] if you prefer the hydronium form)

    • Water autoionization constant: K_w = [\text{H}^+] [\text{OH}^-] = 1.0\times 10^{-14}\;\text{M}^2\;\text{at }25^{\circ}\text{C}

    • Therefore, \text{pH} + \text{pOH} = 14 at 25°C.

  • Connecting concentrations to acidity/basicity:

    • If a solution has a higher [H^+] (e.g., [\text{H}^+] = 10^{-3} \text{ M}, pH = 3), it is more acidic.

    • If another solution has a higher [OH^-] (e.g., [\text{OH}^-] = 10^{-12} \text{ M}, pOH = 12, pH = 2), it is more acidic; however the key is to compare with the corresponding counterpart (hydronium vs hydroxide) and consider pH or pOH accordingly.

  • Practical rule of thumb (from the transcript guidance):

    • Higher hydronium concentration indicates more acidic solutions.

    • Higher hydroxide concentration indicates more basic/alkaline solutions.

    • Neutral solutions have equal hydronium and hydroxide concentrations: [\text{H}^+] = [\text{OH}^-], i.e., pH = pOH = 7 at 25°C.

  • How to decide acidity vs basicity from a given pair of concentrations:

    • If a solution’s hydronium concentration is larger than its hydroxide concentration, the solution is acidic.

    • If a solution’s hydroxide concentration is larger than its hydronium concentration, the solution is basic.

    • If they are equal, the solution is neutral.

  • Quick application reminder: when given numbers in the form 10^{-n}, remember that the larger the exponent (e.g., -3 vs -13), the smaller the concentration; relate this to the pH/pOH values accordingly.

Strong vs Weak Acids and pKa

  • pK_a is a number that indicates acid strength in a numerical form:

    • Smaller pK_a => stronger acid (donates protons more readily).

    • The relationship: pKa = -\log{10} K_a.

  • Common ranges and examples:

    • Seven strong acids (in broad curricula): typically include HCl, HBr, HI, HNO3, H2SO4, HClO4, HClO3 (and sometimes HClO2 depending on the course).

    • A strong acid has a pK_a less than 0 (often < 0 in many charts).

    • Weak acids have pK_a in the range roughly from about 1 to 7.

  • Practical use:

    • When comparing two acids, the one with the smaller pK_a is the stronger acid.

  • Note on teaching videos vs notes:

    • The instructor emphasizes pKa as a numerical way to compare acid strengths, and also notes the strong-acid category (often given as a list of seven strong acids) as a quick reference in buffers questions.

Buffers and Buffered Solutions

  • What buffers do:

    • Buffers resist changes in pH upon additions of small amounts of acid or base.

    • They are essential where reactions require a favorable pH (e.g., physiological pH in blood ~ 7.4).

  • Common buffer composition:

    • Buffers typically consist of a weak acid and its conjugate base, or a weak base and its conjugate acid (buffer pairs).

    • A buffered pair is two species that are conjugates of each other (one hydrogen difference between them).

  • Key buffered pairs (examples):

    • Acetic acid and acetate: \mathrm{CH3COOH}\;\leftrightarrow\; \mathrm{CH3COO^-} (acid ↔ conjugate base)

    • Ammonia and ammonium: \mathrm{NH3}\;\leftrightarrow\; \mathrm{NH4^+}

    • Carbonic acid and bicarbonate: \mathrm{H2CO3}\;\leftrightarrow\; \mathrm{HCO_3^-}

  • Salt involvement in buffers:

    • Buffers can be formed from a weak acid and its conjugate base, or from a weak acid and the salt of that weak acid (the conjugate base).

    • Example: Acetic acid with its salt sodium acetate, \mathrm{CH3COOH} + \mathrm{NaOAc} \rightarrow \, \mathrm{CH3COO^-} + \mathrm{Na^+} + \mathrm{H_2O} (in many buffer formulations, the salt is present as a conjugate base).

    • The conjugate base (e.g., acetate) is the salt form where the hydrogen has been replaced by a metal cation (e.g., Na^+).

  • Why salts sometimes appear in buffers:

    • The salt provides the conjugate base (or conjugate acid) in solution, helping buffer capacity.

    • Not all buffers require a salt; some depend solely on the weak acid and its conjugate base.

  • How to identify buffer candidates (from species given):

    • Look for a weak acid with a conjugate base, or a weak base with a conjugate acid.

    • Ensure the component set does not include a strong acid or a strong base (buffers require weak acid/base components).

    • Some acid-base systems shown in the transcript include acetic acid/acetate, hydrofluoric acid/fluoride, ammonium chloride/ammonia, phosphoric acid/phosphate, carbonic acid/bicarbonate, and acetates with amines.

  • Typical acidic buffers and basic buffers (rules):

    • Acidic buffers: maintain pH < 7; prepared by a weak acid and its conjugate base or the salt of that weak acid.

    • Basic buffers: maintain pH > 7; prepared by a weak base and its conjugate acid or the salt of that weak base.

  • Important cautions from the transcript:

    • Do not include a strong acid or a strong base in a buffer.

    • A buffer’s conjugate pair may appear as a salt; if a salt is present, note that it can be the conjugate of the weak acid/base.

    • In some examples, buffering involves either a weak acid and its conjugate base, or a weak acid with the salt of the weak acid; the latter is equivalent to having the conjugate base present.

  • Examples of buffer species and explanations:

    • Acetic acid/acetic acid salt (sodium acetate) forms an acetate buffer (acidic buffer).

    • Hydrofluoric acid and potassium fluoride form a fluoride buffer (weaker acid with its conjugate base).

    • Ammonia and ammonium chloride form an ammonium buffer (base with conjugate acid).

    • Phosphoric acid and phosphate form a phosphate buffer; carbonic acid and bicarbonate form a bicarbonate buffer.

    • Ethanoic acid and sodium ethanoate illustrate a salt-present buffer going through buffering action.

    • Ethanol (CH3CH2OH) and its conjugate species can be discussed as a buffer in some contexts, though ethanol is a very weak acid and not a common buffering system in water.

  • Practical test-taking tips from the transcript:

    • Draw the structural formula for clarity when solving buffer questions.

    • If a problem presents a salt and an acid/base pair, identify whether the system contains weak acid/base pairs rather than strong acids/bases.

Practical Examples and Common Buffers

  • Acetate buffer:

    • Acid: \mathrm{CH_3COOH}

    • Conjugate base: \mathrm{CH_3COO^-}

    • Salt form: \mathrm{CH_3COONa} (often used to provide the conjugate base in solution)

  • Fluoride buffer:

    • Acid: \mathrm{HF}

    • Conjugate base: \mathrm{F^-}

    • Salt form: \mathrm{KF}

  • Ammonium buffer:

    • Base: \mathrm{NH_3}

    • Conjugate acid: \mathrm{NH_4^+}

    • Salt form: e.g., \mathrm{NH_4Cl}

  • Phosphate buffer:

    • Involves \mathrm{H2PO4^-} / \mathrm{HPO_4^{2-}} pairs (and sometimes Na salts)

  • Carbonate/bicarbonate buffers:

    • Involves \mathrm{H2CO3} / \mathrm{HCO3^-} (and in some cases \mathrm{NaHCO3} salt)

  • Other examples mentioned: ammonia/methlyamine pairs, pyridine/pyridinium, etc., to illustrate conjugate-base pairs in buffers.

  • Non-examples (buffers are not buffers):

    • Strong acids with their conjugates (e.g., HCl with Cl^−) cannot form buffers because the strong acid dominates.

    • Strong bases with their conjugates (e.g., LiOH with Li^+? and related salts) do not form buffers for the same reason.

  • Summary check-list for buffers:

    • Contains a weak acid or weak base and its conjugate pair (base or acid) in solution.

    • May include a salt form for the conjugate partner.

    • No strong acids or strong bases present.

    • The buffer’s pH is determined by the pKa (or pKb) of the weak species involved.

Organic vs Inorganic Chemistry: Quick Distinctions

  • General qualitative differences:

    • Organic compounds typically contain carbon and hydrogen (often with other nonmetals) and are dominated by covalent bonding.

    • Inorganic compounds typically involve elements other than carbon and hydrogen and often feature ionic bonding.

  • Physical properties often differ:

    • Organic: lower melting points and boiling points, often flammable, usually soluble in organic solvents but less soluble in water (solubility tends to decrease with longer carbon chains).

    • Inorganic: higher melting/boiling points, often non-flammable, frequently ionic and highly soluble in water due to ion–water interactions.

  • Solubility definitions:

    • Soluble: substances that mix with water (no visible separate layers).

    • Insoluble: layers or precipitates form (e.g., a solid in a liquid).

  • Role of carbon chains in solubility:

    • Longer hydrocarbon chains reduce water solubility; shorter chains (and the presence of polar functional groups) increase solubility in water.

  • Organic functional groups and naming basics introduced:

    • Alkanes: carbon–carbon single bonds; saturated hydrocarbons. General formula for alkanes: \mathrm{CnH{2n+2}}.

    • Alkanes are named by the prefix corresponding to the number of carbons in the longest continuous chain, followed by -ane (e.g., ethane, propane, butane, pentane, hexane, heptane).

  • Examples of simple alkanes and formulas:

    • Ethane: \mathrm{C2H6}

    • Propane: \mathrm{C3H8}

    • Butane: \mathrm{C4H{10}}

    • Pentane: \mathrm{C5H{12}}

  • Structural representations:

    • Molecular formula shows exact numbers of each type of atom, e.g., \mathrm{C5H{12}}.

    • Condensed structural formulas show how atoms are connected without drawing all bonds: e.g., \mathrm{CH3CH2CH2CH3} for a straight-chain pentane fragment.

    • Skeletal (line-angle) formulas simplify drawings by representing carbon atoms as vertices where lines meet, with hydrogens implied.

    • Expanded structural formulas show all bonds explicitly; condensed formulas are shorthand for sequences of CH groups.

  • Important bonding and structure concepts:

    • Single bonds (C–C or C–H) allow rotation around the bond; double bonds restrict rotation (planarity and static geometry).

    • Carbon valence/Hank-style rules (useful shorthand in class):

    • Hydrogen: 1 bond

    • Oxygen: 2 bonds

    • Nitrogen: 3 bonds

    • Carbon: 4 bonds
      ext{H}
      ightarrow 1, ext{O}
      ightarrow 2, ext{N}
      ightarrow 3, ext{C}
      ightarrow 4

    • Terminal carbons are the ends of a carbon chain; middle carbons are internal.

  • Haloalkanes (halogenated alkanes):

    • Alkanes with halogen substituents (e.g., chlorine, fluorine) still primarily rely on C–C and C–H single bonds and are typically considered organic (though some can form salts or interact ionically in reactions).

  • Polyatomic ions and salts (brief reminder):

    • Some ionic species are polyatomic ions (e.g., \mathrm{SO4^{2-}}, \mathrm{PO4^{3-}}) and are treated as single units in ionic compounds.

  • Practical study tips:

    • In tests, drawing structural formulas helps avoid confusion with condensed or skeletal representations.

    • Know that the same molecule can be written in several formats (molecular, condensed, structural, skeletal), and all can be valid representations.

  • Quick glossary reminders:

    • Salts typically form from metal cations (often group 1 or 2 metals) and anions (like halides or polyatomic ions).

    • A saturated hydrocarbon (alkane) has only single bonds; unsaturated hydrocarbons (alkenes, alkynes) have double or triple bonds respectively.

    • In water, many inorganic compounds dissolve into ions, contributing to conductivity, while many organic compounds do not dissociate significantly.

Summary of Key Formulas and Concepts to Memorize (LaTeX)

  • Acid-base and pH relationships:

    • ext{pH} = -\log_{10} [\text{H}^+]

    • K_w = [\text{H}^+][\text{OH}^-] = 1.0\times 10^{-14}\;\text{M}^2

    • \text{pH} + \text{pOH} = 14 (at 25°C)

  • pKa and acid strength:

    • pKa = -\log{10} K_a

  • Buffers (general):

    • Weak acid and conjugate base: \mathrm{HA} / \mathrm{A^-}

    • Weak base and conjugate acid: \mathrm{B} / \mathrm{BH^+}

    • Acid-base neutralization in buffers (example): \mathrm{CH3COOH} + \mathrm{NaOH} \rightarrow \mathrm{CH3COONa} + \mathrm{H_2O}

  • Alkanes and general formulas:

    • General alkane formula: \mathrm{CnH{2n+2}}

    • Examples:

    • \mathrm{CH4},\; \mathrm{C2H6},\; \mathrm{C3H8},\; \mathrm{C4H{10}},\; \mathrm{C5H_{12}}

  • Bond types and representations:

    • Single bond: one line; double bond: two lines; triple bond: three lines

    • Skeletal formula: vertices/pikes represent carbons; hydrogens implied

Study Tips and Real-World Relevance

  • Blood pH homeostasis is a practical example of buffering (blood ~7.35–7.45).

  • Buffers are essential in biology, chemistry, and environmental science to maintain stable pH for reactions and processes.

  • Understanding buffer composition helps you predict the direction of pH change when acids/bases are added and when salts are involved.

  • For organic chemistry, always practice translating between molecular, condensed, and skeletal representations; this reduces confusion during exams and problem sets.

Note on transcript accuracy: The discussion in the transcript includes helpful intuition but contains a few slips (for example, a 10^{-12} concentration corresponds to pH 12, not pH 2; this is corrected here). The core concepts—buffer pairs, the role of weak acids/bases, conjugates, salts, pKa, and the qualitative distinctions between organic/inorganic chemistry—are accurate and essential for your exam preparation.