Introduction to Acids and Bases

  • Arrhenius Definition:

    • An acid is defined as a substance that contains a hydrogen atom and dissolves in water to form a hydrogen ion, $H^+$.

    • A base is defined as a substance that contains hydroxide ($−OH$) and dissolves in water to form $−OH$.

    • Example:

    • Hydrochloric Acid (HCl): $HCl(g)
      ightarrow H^+(aq) + Cl^−(aq)$

    • Sodium Hydroxide (NaOH): $NaOH(s)
      ightarrow Na^+(aq) + −OH(aq)$

Limitations of the Arrhenius Definition

  • The Arrhenius definition accurately predicts the behavior of many acids and bases.

  • However, it is limited and can be inaccurate in some situations.

    • Example: The hydrogen ion $H^+$ does not exist independently in water; it reacts with water to form the hydronium ion $H_3O^+$.

    • This is represented:

    • $H^+(aq) + H2O(l) ightarrow H3O^+(aq)$

    • It is important to note that:

    • Hydrogen Ion: Does not exist in solution.

    • Hydronium Ion: Is actually present in aqueous solution.

Brønsted–Lowry Definition

  • The Brønsted–Lowry definition is a more widely accepted concept:

    • A Brønsted–Lowry acid is a proton ($H^+$) donor.

    • A Brønsted–Lowry base is a proton ($H^+$) acceptor.

    • Example of this process:

    • $H3O^+(aq) + Cl^− ightarrow HCl(g) + H2O(l)$

    • Here, $HCl$ is a Brønsted–Lowry acid because it donates a proton to water, which acts as a Brønsted–Lowry base by accepting a proton from $HCl$.

Common Brønsted–Lowry Acids

  • A Brønsted–Lowry acid must contain a hydrogen atom.

  • Common examples of Brønsted–Lowry acids (HA) include:

    • Hydrochloric acid (HCl)

    • Hydrobromic acid (HBr)

    • Acetic acid (CH₃COOH)

    • Sulfuric acid (H₂SO₄)

    • Nitric acid (HNO₃)

Brønsted–Lowry Bases

  • A Brønsted–Lowry base is a proton acceptor capable of forming a bond to a proton.

  • A base must contain a lone pair of electrons that can be used to form a bond to a proton.

  • Example Illustration:

    • $N ext{ with lone pair } + H_2O(l)
      ightarrow N-H + −OH(aq)$

    • In this reaction, the electron pair forms a new bond with a $H^+$ from water, illustrating a Brønsted–Lowry base structure.

Common Brønsted–Lowry Bases

  • Common Brønsted–Lowry bases (B) include:

    • Sodium Hydroxide (NaOH)

    • Potassium Hydroxide (KOH)

    • Magnesium Hydroxide (Mg(OH)₂)

    • Calcium Hydroxide (Ca(OH)₂)

    • Water (H₂O)

    • Ammonia (NH₃)

  • These bases are neutral compounds that exhibit basicity primarily due to their capacity to induce a reaction involving $−OH$.

Proton Transfer in Acid-Base Reactions

  • Conjugate Acid-Base Pairs:

    • The product formed by the loss of a proton from an acid is referred to as its conjugate base.

    • The product formed by the gain of a proton by a base is called its conjugate acid.

    • Reaction Example Representation:

    • $HA + B
      ightleftharpoons A^- + B^+$

    • Where $HA$ is the acid, $B$ is the base, $A^-$ is the conjugate base, and $B^+$ is the conjugate acid.

Examples of Conjugate Acid-Base Pairs

  • Example of conjugate pairs:

    • $HBr$ and $Br^−$ as a conjugate acid-base pair.

    • $H2O$ and $H3O^+$ as a conjugate acid-base pair.

  • Important aspect: The net charge must remain equal on both sides of the reaction equation.

Charge Changes in Proton Transfer

  • When a species gains a proton ($H^+$), it gains a charge of +1.

    • Example:

    • $H2O (neutral) + H^+ ightarrow H3O^+ (+1 charge)$

  • Conversely, when a species loses a proton ($H^+$), it gains a charge of -1.

    • Example:

    • $HBr (neutral) - H^+
      ightarrow Br^− (-1 charge)$

Amphoteric Compounds

  • Amphoteric Compound: A compound capable of acting as either an acid or a base, containing both a hydrogen atom and a lone pair of electrons.

  • Example of behavior of $H_2O$:

    • As a Base:

    • $H2O + H^+ ightarrow H3O^+$ (gaining $H^+$ forming conjugate acid)

    • As an Acid:

    • $H_2O - H^+
      ightarrow OH^- $ (losing $H^+$ forming conjugate base)

Identifying Acid-Base Components

  • Practice Identifying:

    • What is the conjugate acid of $NO_3^-$?

    • What is the conjugate base of $HCO_3^-$?

    • Example Reaction:

    • $CH3COOH (l) + NH3 (g)
      ightarrow CH3COO^− (aq) + NH4^+ (aq)$

  • Identifying components:

    • Acid: $CH_3COOH$

    • Base: $NH_3$

    • Conjugate Acid: $NH_4^+$

    • Conjugate Base: $CH_3COO^−$

Acid and Base Strength

  • Dissociation Dynamics:

    • When a covalent acid dissolves in water, the proton transfer that forms $H_3O^+$ is termed dissociation.

  • Strong vs. Weak Acids:

    • Strong Acids:

    • An example includes $HI, HBr, HCl, H2SO4, HNO_3$; they fully dissociate in water (100%) forming ions.

    • Representation:

      • $H3O^{+}(aq) + Cl^{-}(aq) ightarrow HCl(g) + H2O(l)$

    • Weak Acids:

    • Weak acids dissociate into ions only partially.

    • Examples: $H3PO4, HF, H2CO3, HCN$.

    • Representation:

      • $H3O^{+}(aq) + CH3COO^{-}(aq)
        ightleftharpoons CH3COOH(l) + H2O(l)$

Observing Strong vs. Weak Acids

  • Comparison:

    • A strong acid like $HCl$ is completely dissociated into $H_3O^{+}(aq)$ and $Cl^{-}(aq)$.

    • A weak acid like $CH_3COOH$ contains mostly unreacted acid.

Base Characteristics

  • Strong Bases:

    • Strong bases like $NaOH$ fully dissociate into ions when dissolved in water.

    • Example:

      • $NaOH(s)
        ightarrow Na^{+}(aq) + −OH(aq)$

    • Common strong bases include $NaOH$ and $KOH$.

  • Weak Bases:

    • Weak bases dissociate into ions only partially in water.

    • Example:

      • $NH4^{+}(aq) + NH3(g) + H_2O(l)
        ightleftharpoons −OH(aq)$

    • Weak bases are typically found as neutral species.

Strong and Weak Acid-Base Relationships

  • A strong acid readily donates protons thereby forming a weak conjugate base. (e.g., $HCl
    ightarrow Cl^{-}$)

  • A strong base readily accepts protons forming a weak conjugate acid (e.g., $−OH
    ightarrow H_2O$).

Water as a Supramolecular Solvent

  • Both acidic and basic traits:

    • Water can behave as both a Brønsted–Lowry acid and a Brønsted–Lowry base, allowing it to undergo self-ionization.

  • Example Reaction:

    • $2H2O (l) ightleftharpoons H3O^{+} (aq) + OH^{-} (aq)$

Dissociation in Pure Water

  • Pure water has an extremely low concentration of ions: $H_3O^+$ and $−OH$.

  • Both ions are formed equally in each reaction:

    • Concentrations of $[H_3O^+]$ and $[−OH]$ are equal in pure water.

  • The Ion-Product Constant for Water, $Kw$, can be calculated by multiplying $[H3O^+]$ and $[−OH]$:

    • K<em>w=[H</em>3O+][−OH]K<em>w = [H</em>3O^+][−OH]

  • At 25 °C, [H3O+]=[−OH]=1.0imes10−7extM[H_3O^+] = [−OH] = 1.0 imes 10^{-7} ext{ M}.

Ion-Product Constant Calculation

  • Calculating $K_w$:

    • Substituting the concentrations yields:

    • Kw=(1.0imes10−7)imes(1.0imes10−7)K_w = (1.0 imes 10^{-7}) imes (1.0 imes 10^{-7})

    • Kw=1.0imes10−14K_w = 1.0 imes 10^{-14}

    • The value of $K_w$ remains a constant $1.0 imes 10^{-14}$ in all aqueous solutions at 25 °C.

pH Scale and Calculation

  • Calculating pH:

    • The formula to determine pH is:

    • pH=−extlog[H3O+]pH = - ext{log} [H_3O^+]

  • Descriptions of Solution Type:

    • Acidic Solution: $pH < 7 ightarrow [H_3O^+] > 1.0 imes 10^{-7}$

    • Basic Solution: $pH > 7
      ightarrow [H_3O^+] < 1.0 imes 10^{-7}$

    • Neutral Solution: $pH = 7
      ightarrow [H_3O^+] = 1.0 imes 10^{-7}$

pH Calculation Examples

  • Example: If [H3O+]=1.0imes10−5extM[H_3O^+] = 1.0 imes 10^{-5} ext{ M}, what is its pH?

  • Example: If the pH of a solution is 8.50, calculate [H3O+][H_3O^+].

Human Body and pH

  • The pH of body fluids varies significantly:

    • Saliva: pH 5.8 - 7.1

    • Esophagus: pH 7.4

    • Stomach: pH 1.6 - 1.8

    • Small Intestine: pH 8.5

    • Large Intestine: pH 5.0 - 7.0

    • Urine: pH 4.6 - 8.0

    • Blood: pH 7.4

Neutralization Reactions

  • Neutralization Reaction: An acid-base reaction yielding salt and water.

    • General Reaction Form:

    • HA(aq)+MOH(aq)<br>ightarrowH2O(l)+MA(aq)HA(aq) + MOH(aq) <br>ightarrow H_2O(l) + MA(aq)

    • In the reaction:

    • The acid $HA$ donates a proton ($H^+$) to $−OH$, forming water ($H_2O$).

    • The anion $A^−$ from the acid combines with the cation $M^+$ from the base to form salt $MA$.

Balanced Equation example

  • Example Question:

    • Write a balanced equation for the reaction of HNO₃ with NaOH.

    • Write a balanced equation for the reaction of Mg(OH)₂ with HCl.

Net Ionic Equations

  • Net Ionic Equation: Focuses exclusively on the species involved in a reaction.

  • For example, in the reaction:

    • HCl(aq)+NaOH(aq)<br>ightarrowH2O(l)+NaCl(aq)HCl(aq) + NaOH(aq) <br>ightarrow H_2O(l) + NaCl(aq)

    • Written in Ionic Form:

    • H+(aq)+Cl−(aq)+Na+(aq)+−OH(aq)<br>ightarrowH2O(l)+Na+(aq)+Cl−(aq)H^+(aq) + Cl^−(aq) + Na^+(aq) + −OH(aq) <br>ightarrow H_2O(l) + Na^+(aq) + Cl^−(aq)

    • Removing Spectator Ions (Na+ and Cl−):

    • H+(aq)+−OH(aq)<br>ightarrowH2O(l)H^+(aq) + −OH(aq) <br>ightarrow H_2O(l)

Acid-Base Reactions with Carbonates

  • Reactions involving bicarbonate bases:

    • $HCO₃^−$ reacts with one proton to form carbonic acid:

    • H+(aq)+HCO3−(aq)<br>ightarrowH2CO3(aq)H^+(aq) + HCO₃^−(aq) <br>ightarrow H_2CO₃(aq)

    • Carbonic acid then decomposes into:

    • H<em>2CO3ightarrowH</em>2O(l)+CO2(g)H<em>2CO₃ ightarrow H</em>2O(l) + CO_2(g)

  • Example:

    • HCl(aq)+NaHCO3(aq)<br>ightarrowH<em>2O(l)+CO</em>2(g)+NaCl(aq)+H2CO3(aq)HCl(aq) + NaHCO₃(aq) <br>ightarrow H<em>2O(l) + CO</em>2(g) + NaCl(aq) + H_2CO₃(aq)

Reactions with Carbonates

  • Reactions involving carbonate bases:

    • CO32−+2H+<br>ightarrowH2CO3CO₃^{2-} + 2H^+ <br>ightarrow H_2CO₃

    • Reaction Example:

    • 2HCl(aq)+CaCO3(aq)<br>ightarrowH<em>2O(l)+CO</em>2(g)+2CaCl2(aq)+H2CO3(aq)2HCl(aq) + CaCO₃(aq) <br>ightarrow H<em>2O(l) + CO</em>2(g) + 2CaCl₂(aq) + H_2CO₃(aq)

Titration

  • Titration: A technique used to determine the concentration of an acid or base in a solution.

    • When determining acid concentration, a base of known concentration is slowly added until neutralization occurs.

    • At neutralization, the # of moles of acid equals the # of moles of base.

    • This end point indicates equivalence.

Determining Molarity in Titration

  • To find the unknown molarity of an acid from titration data:

    • Calculate moles of base used.

    • Use molarity of acid solution with a mole-mole conversion factor.

    • Determine final concentration:

    • extMolesofacid=extVolumeofbaseimesextMolarityofbaseext{Moles of acid} = ext{Volume of base} imes ext{Molarity of base}

Example Titration Calculation

  • Question: What is the molarity of an HCl solution if 22.5 mL of a 0.100 M NaOH solution are needed to titrate a 25.0 mL sample of the acid?

Buffers

  • A Buffer is a solution that resists changes in pH upon the addition of acids or bases.

    • Most buffers consist of roughly equal parts:

    • A weak acid and its conjugate base in salt form.

  • Functionality:

    • When a base ($−OH$) is added, it reacts with the weak acid preventing drastic pH change.

    • When an acid ($H_3O^+$) is added, it reacts with the conjugate base, again minimizing pH variation.

General Buffer Characteristics

  • Considered in equilibrium reactions:

    • $CH3COOH(aq) + H2O(l) H3O^+(aq) + CH3COO^−(aq)$

  • If extra acid is introduced:

    • It reacts with the conjugate base, driving the equilibrium backward, minimizing pH alterations.

Buffer Actions with Added Base

  • Example Reaction in Buffer:

    • $CH3COOH(aq) + −OH(aq) ightarrow H2O(l) + CH_3COO^−(aq)$

    • Additional base drives the equilibrium forward, thereby minimizing pH change.

Buffers in the Human Body

  • Normal Blood pH: Ranges between 7.35 and 7.45.

  • Key Buffer System: Carbonic Acid/Bicarbonate System (CO₂(g) + H₂O(l) ightleftharpoons H₂CO₃(aq) ightleftharpoons H₃O^+(aq) + HCO₃^−(aq) ).

    • CO₂ is a product of metabolism influencing blood pH.

Respiratory Impact on Blood pH

  • Respiratory Acidosis: Occurs when CO₂ is retained in the body due to lung issues.

    • Increased $CO₂$ drives the reverse reaction to the right, boosting $H_3O^+$ levels and lowering pH.

  • Respiratory Alkalosis: Occurs through hyperventilation, leading to a decrease in $CO₂$.

    • Decreasing $CO₂$ shifts the reaction to the left, lowering $H_3O^+$ and increasing pH.