Chapter 17: Acids and Bases

Properties of Acids

• General Characteristics:

  • Acids are commonly used in laboratory and industrial settings.

  • Example of laboratory use: Dissolving metals in soil using acids like nitric acid.

• Common Acids

  • Acetic Acid: Key component of vinegar, contributes sour taste.

  • Hydrochloric Acid:

  • Used in metal cleaning and cleaning products.

  • Major acid in the human stomach, responsible for acid reflux and heartburn.

  • Sulfuric Acid: Used in fertilizer production and batteries.

  • Nitric Acid: Widely used in fertilizers.

  • Citric Acid: Found in fruits, contributes to tartness.

  • Carbonic Acid: Present in carbonated drinks; derived from carbon dioxide in water.

  • Hydrofluoric Acid: Extremely corrosive, can etch glass, dangerous to handle.

  • Phosphoric Acid: Commonly used in fertilizers.

  • Ammonium ($NH_4^{+}$ ): It acts as an acid by donating a proton ($H^{+}$ ) to a base, reverting to ammonia ($NH_3$ ).

• Naming Acids:

  • Distinctions based on suffixes:

  • Acids ending in -ate lead to the suffix -ic in the acid name.

  • Acids ending in -ite lead to the suffix -ous.

  • Example: Sulfate → Sulfuric Acid; Sulfite → Sulfurous Acid.

  • Monoprotic: has one H

    Diprotic: has two H

    Triprotic: has three H

Properties of Bases

• General Characteristics:

  • Bases typically have a bitter taste (e.g., coffee, chocolate) and feel slippery (due to soap-like consistency when mixed with water).

  • Bases are present in cleaning products (e.g., drain cleaners).

• Common Bases:

  • Sodium Hydroxide: Frequently used in laboratories.

  • Potassium Hydroxide: Similar applications as sodium hydroxide.

  • Sodium Bicarbonate (Baking Soda): Used in cooking, neutralizes stomach acid.

  • Sodium Carbonate: Used to soften water and in glass production.

  • Ammonia (NH₃): Acts as a base but does not contain hydroxyl ions in its natural form.

Definitions of Acids and Bases

Three definitions: Arrhenius, Bronsted-Lowry, Lewis.

No individual definition of acids and bases is correct, each is useful in a given instance

• Arrhenius Definition:

  • Acid: Substance that produces hydrogen ions (H⁺) in aqueous solution.

  • Base: Substance that produces hydroxide ions (OH⁻) in aqueous solution.

  • Example: Hydrochloric acid (HCl) dissociates in water to form H⁺ and Cl⁻.

  • Acids and bases combine to form water, thereby neutralizing each other

  • acid + base → salt + water

  • Example: HCl (𝑎𝑞) + NaOH (𝑎𝑞) → NaCl (𝑎𝑞) + H$_2$O (𝑙)

• Bronsted-Lowry Definition:

  • Acid: Proton ($H^{+}$ion) donor.

  • Base: Proton ($H^{+}$ion) acceptor.

  • Example: In the reaction of ammonia with water, ammonia accepts a proton to become ammonium (NH₄⁺).

• Lewis Definition:

  • Lewis Acid: Electron pair acceptor.

  • Lewis Base: Electron pair donor.

  • This definition expands the scope of acids and bases beyond just protons.

Arrhenius Acid-Base Reactions

• Neutralization Reaction:

  • Reacting an acid with a base yields water and a salt.

  • Example: The reaction between hydrochloric acid (HCl) and sodium hydroxide (NaOH) produces sodium chloride (table salt) and water (H₂O).

Hydronium Ion

The $H^{+}$ ions produced by the acid are so reactive that they cannot exist in water.

$H^{+}+H_2O=H_3O^{+}$

• Formation of the hydronium ion (H₃O⁺) from hydrogen ions and water:

  • When protons (H⁺) combine with water molecules, hydronium ions are formed, indicating an acidic solution.

Chemists often use H+ (aq) and H3O+ (aq) interchangeably to mean the same thing – an H+ ion that has dissolved in water

Bronsted-Lowry Definition

Acid: Proton (H+ ion) donor

Base: Proton (H+ ion) acceptor

• All Arrhenius acids and bases are acids/bases under the Brønsted-Lowry definition. However, some B-L acids and bases are not Arrhenius acids/bases.

$NH_3\left(aq\right)+H_2O\left(l\right)\rightleftharpoons NH_4^{+}\left(aq\right)+OH^{-}\left(aq\right)$

NH3 does not inherently contain OH$^{-}$ ions but it still produces OH$^{-}$ ions in solutions.

According the B-L definition, $NH_3$ is a base because it accepts a proton from water

Amphoteric Substances

According to the B-L definition, substances (such as H2O) can act as acids or bases

H2O as donor (acid): $NH_3\left(aq\right)+H_2O\left(l\right)\rightleftharpoons NH_4^{+}\left(aq\right)+OH^{-}\left(aq\right)$

H2O as acceptor (base): $NH_4\left(aq\right)+H_2O\left(l\right)\rightleftharpoons NH_3\left(aq\right)+H_3O\left(aq\right)$

Acid-Base Pairs and Conjugates

• Conjugate Acids and Bases:

  • A conjugate acid is the species formed after a base gains a proton.

  • A conjugate base is the species that remains after an acid has donated a proton.

  • Example: Ammonium (NH₄⁺) is the conjugate acid of ammonia (NH₃).

Strong vs. Weak Acids and Bases

The dissolved ions act as charge carriers, thereby conducting electricity

• Substances that completely dissociate into ions in water are strong electrolytes

• Compounds that do not dissociate into ions in water are nonelectrolytes

• Strong Acids: Completely dissociate in water and are strong electrolytes ($\rightarrow$)

  • Examples: Hydrochloric acid (HCl), hydrobromic acid (HBr), sulfuric acid (H₂SO₄, diprotic).

  • All hydrohalic (derived from halogens) acids are strong acids that completely ionize in water except hydrofluoric acid (HF)

• Weak Acids: Do not completely dissociate in water and are weak electrolytes (⇌).

  • Examples: Hydrofluoric acid (HF), acetic acid (CH₃COOH), carbonic acid (H₂CO₃, diprotic).

The degree to which an acid is strong or weak depends on the attraction between the anion of the original acid (the conjugate base) and the H+, relative to the attractions of these ions to water.

$HA\left(aq\right)+H_2O\left(l\right)\rightleftharpoons H_3O^{+}\left(aq\right)+A^{-}\left(aq\right)$

• If the attraction between H+ and A is weak $\rightarrow$ the forward direction is favored and the acid is strong

• If the attraction between H+ and A is strong $\larr$ the reverse direction is favored and the acid is weak

Dissociation Constants ($K_{a}$ value)

• Ka (Acid Dissociation Constant):

  • A measure of the strength of an acid in solution.

  • Higher Ka indicates a stronger acid.

  $HA\left(aq\right)+H_2O\left(l\right)\rightleftharpoons H_3O^{+}\left(aq\right)+A^{-}\left(aq\right)$

  • The equilibrium expression for dissociation is given by: $K_{a}=\frac{[H_3O^{+}][A^{-}]}{[HA]}$ , where [HA] is the concentration of the undissociated acid.

  • [$H_2O$ ] is not included in calculations as it is a pure liquid.

Amphoteric Water ($K_{W}$)

• Water is amphoteric: can act as both an acid and a base.

  • Acts as a base when mixed with an acid (hydrofluoric acid in this case).

  • Produces hydronium ions (H₃O⁺) when acting as a base.

  • Acts as an acid when mixed with a base, producing hydroxide ions (OH⁻).

• Water can autoionize, producing both hydronium and hydroxide ions on its own.

  • Reaction: $H_2O\left(l\right)+H_2O\left(l\right)\rightarrow H_3O^{+}\left(aq\right)+OH^{-}\left(aq\right)$

• Identifying the ion product/dissociation constant of water ($K_{W}$ ) at 25°C: no division by reactants as it is just pure liquid water

$K_{W}=\left\lbrack H_3O^{+}\right\rbrack\left\lbrack OH^{-}\right\rbrack$ or $\left\lbrack H^{+}\right\rbrack\left\lbrack OH^{-}\right\rbrack$
$K_{w}=1.0\cdot10^{-14}$

• At neutral pH:

  • Concentration of H₃O⁺ and OH⁻ is equal:

  • $K_{W}=\left\lbrack H_3O^{+}\right\rbrack^2$ or $K_{W}=\left\lbrack OH^{-}\right\rbrack^2$

  • $[H_3O^{+}]=[OH^{-}]=\sqrt{K_{W}}=1.0\cdot10^{-7}$

• Importance of $K_{W}$ : Provides baseline for acid-base reactions.

  • At neutral pH, both H₃O⁺ and OH⁻ equal each other.

Calculating Hydroxide Ion Concentration using $K_{W}$

• For a given $[H_3O^+]$, calculate $[OH^-]$ using:

  • $Kw=[H3O^{+}][OH^{-}]=1.0\cdot10^{-14}$

  • If $[H_3O^+] = 7.5 \times 10^{-5}$ :

  • $[OH^{-}]=\frac{Kw}{[H_3O^{+}]}$

  • Example Calculation: $\left\lbrack OH^{-}\right\rbrack=\frac{10^{-14}}{7.5 \times10^{-5}}\rightarrow1.33\times10^{-10}$

• Identify if acidic, basic, or neutral:

  • If [H_3O^+] > [OH^-], solution is acidic.

  • If [H_3O^+] < [OH^-], solution is basic.

pH (hydronium concentration)

• pH is defined as the negative log of hydronium concentration:
  $pH = -\log[H_3O^+]$

• Example: If $[H_3O^+] = 1 \times 10^{-3}$ , then:

  • $pH = -\log(1 \times 10^{-3}) = 3$

Sig Figs and pH

• Sig figs are essential for pH reporting.

• If $[H_3O^+]$ has:

  • Two significant figures: pH should have two decimal places.

  • Three significant figures: pH should have three decimal places.

Characteristics of pH scale

• Less than 7 (pH < 7): acidic

• More than 7 (pH > 7): basic

• Exactly 7 (pH = 7): neutral

• Each unit change in pH represents a tenfold change in acidity:

  • Example: pH 2 (Lime) is 10x more acidic than pH 3 (Plum).

  • pH changes from 2 to 4 represent a 100x change in acidity.

Example Problem 1

• Determine pH from given hydronium concentration $[H_3O^{+}]=1.8\times10^{-4}$

$pH=-\log\left(1.8\cdot10^{-4}\right)$ = 3.74 < 7 = acidic

Example Problem 2

• Calculate the pH if $[OH^{-}]=1.3\times10^{-2}$ .

First solve for $\left\lbrack H_3O^{+}\right\rbrack$ concentration

$\left\lbrack H_3O^{+}\right\rbrack\left\lbrack OH^{-}\right\rbrack=K_{W}$

$\left\lbrack H_3O^{+}\right\rbrack\left(1.3\cdot10^{-2}\right)=1.0\cdot10^{-14}$

H$\left\lbrack H_3O^{+}\right\rbrack=\frac{1.0\cdot10^{-14}}{1.3\cdot10^{-2}}=7.7\cdot10^{-13}M$

Then use pH formula

$pH=-\log\left\lbrack7.7\cdot10^{-13}\right\rbrack=-\left(-12.11\right)=12.11$ > 7 = basic

Example Problem 3

• Calculate $\left\lbrack H_3O^{+}\right\rbrack$ concentration for a solution with a pH of 4.80

$4.80=-\log\left\lbrack H_3O^{+}\right\rbrack$

$-4.80=\log\left\lbrack H_3O^{+}\right\rbrack$

using this:$10^{\log x}=x$

$10^{-4.80}=10^{\log\left\lbrack H_3O^{+}\right\rbrack}$

$\left\lbrack H_3O^{+}\right\rbrack=10^{-4.80}=1.6\cdot10^{-5}$ $

pOH (hydroxide concentration)

• $pOH = -\log[OH^-]$

  • When pOH is less than 7: solution is acidic.

  • When pOH is greater than 7: solution is basic.

• pH + pOH = 14 (at 25°C)

Strength of Acids ($_{pK}$)

• Acid strength is determined by dissociation constant (Kₐ).

• Large Kₐ = strong acid.

• pKₐ is calculated as:
  $pKa = -\log(Ka)$

• $K_{a}=10_{}^{-pK_{a}}$

• Stronger acids have larger $K_{a}$ but smaller $pK_{a}$ values.

• Bases have the same: $K_{b}$ and $pK_{b}$

Two Sources of Hydronium

• Hydronium can come from:

  1. Dissociation of the acid

  2. Autoionization of water.

• $H_2O\left(l\right)+H_2O\left(l\right)\rightarrow H_3O^{+}+OH^{-}$

Except in extremely dilute acid solutions (< 10-5 M), the autoionization of water contributes a negligibly small amount of H3O+ compared to the ionization of the acid

The addition of the H3O+ from the acid actually shifts the autoionization reaction to the left, producing even less H3O+

Strong and Weak Acids

• Strong acids fully dissociate in water (e.g. HCl).

  • The concentration of H₃O⁺ equals the concentration of the added acid.

• Weak acids only partially dissociate (e.g. HF); their pH must be calculated using equilibrium expressions.

Strong Acids: 100% dissociation

• Hydrochloric Acid (HCl)

• Hydrobromic Acid (HBr)

• Hydroiodic Acid (HI)

• Upper of polyatomic series

Weak Acids: <5%

• Hydrofluoric Acid (HF)

• Acetic Acid (CH₃COOH)

• Citric Acid (C₆H₈O₇)

• Lactic Acid (C₃H₆O₃)

• Formic Acid (CH$_2$ O$_2$)

• Carbonic Acid (H$_2$ CO$_3$)

• Hydrogen Cyanide (HCN)

• Lower of polyatomic series

Rules

Large organic compounds are going to be weak acids

In binary acids, all the halogens are strong acids except HF

In oxyacid series, the number of oxygens determines the acid strength

• Per-ic and ic acids are strong ($HClO_4$ /$HClO_3$ )

• Hypo-ous and ous acids are weak ($HClO_2$ /$HClO$ )

Strong Acids

• We simply assume:

  [H+] = initial acid concentration

So in practice:

• Strong acids are not analyzed using $K_{a}$ equilibrium calculations.

• They are treated as 100% dissociated.

$0.10M$ HCl → $\left\lbrack H_3O^{+}\right\rbrack=0.10M$ → $pH=-\log\left(0.10\right)=1.00$

Weak Acid Calculations

The concentration of H3O+ in a weak acid is not equal to the concentration of the weak acid - < 5% dissociation.

To solve for concentration:

• Set up ICE table:

  • Initial concentrations, equilibrium changes, and final concentrations.

  • Apply the equilibrium constant expression ($K_{a}$) and solve.

  • Use the ‘x is small” assumption if valid

  Valid if $\frac{x_{approximate}}{initial}\cdot100$ < 5%

• If using the assumption that x is small is invalid, apply the quadratic formula

Solving for concentration of $H_3O^{+}$

$HA\left(aq\right)+H_2O\left(l\right)\rightleftarrows H_3O^{+}\left(aq\right)+A^{-}\left(aq\right)$ , $\left\lbrack HA\right\rbrack=0.10M$

$H_2O\left(l\right)$ is a pure liquid and therefore not included in the $K_{a}$ constant calculation

$K_{a}=\frac{\left\lbrack H_3O^{+}\right\rbrack\left\lbrack A^{-}\right\rbrack}{\left\lbrack HA\right\rbrack}=\frac{\left(x\right)\left(x\right)}{0.10-x}=\frac{x^2}{0.10-x}$

Example: Solving for the concentration of H3O+ of a weak acid solution (x is small)

Example: Solving for the concentration of H3O+ of a weak acid solution (x is not small)

Solving for $K_{a}$ from pH of a weak acid

Using given pH, the concentration of hydronium ions can be calculated

Then using an ICE table, $K_{a}$ can be calculated

Percent Ionization

• Definition: The percentage of the initial weak acid that ionizes in solution:

  • Percent Ionization: $\frac{\left\lbrack H_3O^{+}\right\rbrack}{\left\lbrack HA\right\rbrack}\cdot100$

  • Where $[{HA}]$ is the initial concentration of the weak acid.

• Characteristics: Most weak acids ionize between 1%-5%, indicating their weak nature because only a small fraction completely dissociates into ions.

Percent ionization decreases as the acid concentration increases.

$HA\left(aq\right)+H_2O\left(l\right)\rightleftarrows H_3O^{+}+A^{-}\left(aq\right)$

• If [HA] increases (acid becomes more concentrated- removal of water) it shifts the reaction to the left and less acid ionizes

• If [HA] decreases (acid becomes more dilute- adding more water) it shifts the reaction to the right and more acid ionizes

Acid Mixture: Strong Acid + Weak Acid

Three sources of H3O+:

• from strong acid

• from weak acid

• from auto ionization of water

The formation of H3O+ by the strong acid suppresses the formation of additional H3O+ formed by the auto ionization of water and USUALLY by the weak acid. However, this needs to be checked:

Example

0.10 M $HCl$ + 0.10 M $HCHO_2$ (formic acid)

• From auto ionization of water: [H3O+ ] ≈ 0.00 (can assume negligible)

• From Strong Acid: [H3O+] = 0.10 M since HCl is a strong acid

• For weak acid: need to solve for [H3O+] or [CHO2- ] (same value at equilibrium)

Weak Acid Equation: $HCHO_2+H_2O\left(l\right)\rightleftarrows H_3O^{+}\left(aq\right)+CHO_2^{-}\left(aq\right)$

• $K_{a}=1.8\cdot10^{-4}$

Since the $HCl$ is a strong acid and already dissociated 100% of itself into $H_3O^{+}$ , the intial$\left\lbrack H_3O^{+}]\right.$ is 0.10 M, along with the given initial $\left\lbrack HCHO_2\right\rbrack$

Acid Mixture: Weak Acid + Weak Acid

Three sources of H3O+:

• from strong acid

• from weak acid

• from auto ionization of water

To calculate the for a mixture of two weak acids:

Compare the values and perform the standard weak acid calculation using only the weak acid with the larger $K_{a}$

• The stronger of the weak acids acts as the dominant source of ions; its dissociation creates a Common Ion Effect that suppresses BOTH the ionization of the weaker acid and water auto ionization

• The contribution from the second acid becomes so small that it does not significantly change the total concentration or the final pH.

Example

0.300 M $HF$ + 0.100 M $HClO$

• From auto ionization of water: [H3O+ ] ≈ 0.00 (can assume negligible)

• Strongest weak acid HF ($K_{a}=6.8\cdot10^{-4}$ )

Stronger Weak Acid Equation: $HF\left(aq\right)+H_2O\left(l\right)\rightleftarrows H_3O^{+}\left(aq\right)+F^{-}\left(aq\right)$

• Remember pure liquids are always excluded

Strong and Weak Bases

• Strong Bases: Bases that completely dissociate in solution, producing hydroxide ($OH^{-}$ ) ions directly:

  • Ex: 1.0 M $NaOH$$\rightarrow$ [$OH^{-}$ ] = 1.0 M and [$Na$ ] = 1.0 M

  • Strong bases are typically alkali and alkaline earth metals.

• Weak Bases: Do not completely dissociate; must calculate hydroxide concentration similarly to weak acids using ICE tables:

• Strong bases ‘contain’ OH- and dissociate in water

• Weak bases ‘produce’ OH- by accepting a proton from water to form OH-

  • Ex: Ammonia (base) $\rightarrow$ Ammonium (conjugate acid)$NH_3\left(aq\right)+H_2O\left(l\right)\rightleftharpoons NH_4^{+}\left(aq\right)+OH^{-}\left(aq\right)$

Weak bases such as ammonia contain an amine group with a N with a lone pair that acts as the proton acceptor.

Amines are generally weak bases because the nitrogen lone pair, while available for protonation, is not extremely eager to accept a proton, especially when stabilized by nearby atoms or resonance

$K_{b}$ - Base Ionization Constant

$Base\left(aq\right)+H_2O\left(l\right)\rightleftharpoons BaseH^{+}\left(aq\right)+OH^{-}\left(aq\right)$

$K_{b}=\frac{\left\lbrack BaseH^{+}\right\rbrack\left\lbrack OH^{-}\right\rbrack}{\left\lbrack Base]\right\rbrack}$

• The smaller the $K_{b}$, the weaker the base

• $pK_{b}=$ -log ($K_{b}$)

Strong Base Calculation

• For strong bases, the ionization is assumed to be complete, so the concentration of the base equals the concentration of the hydroxide ions produced (with stoichiometry)

• To find hydronium ion concentration, we need to use the $K_{W}$ formula

$\left\lbrack H_3O^{+}\right\rbrack\left\lbrack OH^{-}\right\rbrack=K_{W}\left(1.0\cdot10^{-14}\right)$

Then we can use the pH formula:

pH = -log ($H_3O^{+}$)

Weak Base Calculation

Anion Behavior (Conjugate Bases)

• Conjugate bases can influence solution properties. The strength of a conjugate base depends on the strength of its parent acid:

  • From Strong Acids: Their conjugate bases are neutral (do not affect pH).

  • From Weak Acids: Their conjugate bases can act as weak bases (increase pH).

Example:

$HF(aq)+H_2O(l)\longrightarrow{}H_3O+(aq)+F^{-}(aq)$

acid, base, conjugate acid, conjugate base

$F^{-}$ (conjugate base) has an affinity for $H^{+}$ ions $\rightarrow$ removing $H^{+}$ from water molecules

$F^{-}(aq)+H_2O(l)\rightleftarrows OH^{-}(aq)+HF(aq)$

Weaker the parent acid = stronger the conjugate base

Solving for pH of a Weak Acid Anion Solution

1. Find Kb: $Kb=\frac{K_{W}}{K_{a}}$

2. Solve for x using ICE table concept: use approximation if valid, or quadratic if needed.

3. Find [H₃O⁺]: $[H3O^{+}]=\frac{K_{W}}{[OH^{-}]}$

4. pH = −log[H₃O⁺]

Cation Behavior (Conjugate Acids)

• Conjugate acids can influence solution properties. The strength of a conjugate acid depends on the strength of its parent base:

  • From Strong Bases: Their conjugate bases are neutral (do not affect pH).

  • From Weak Bases: Their conjugate acids can act as weak acids (decrease pH).

Example:

$NH_3\left(aq\right)+H_2O\left(l\right)\rightleftharpoons NH_4^{+}\left(aq\right)+OH^{-}$

acid, base, conjugate acid, conjugate base

$NH_{^{}4}^{^{+}}$ (conjugate acid) donates $H^{+}$ ions to water $\rightarrow$ turning it back into $NH_3$ and producing $H_3O^{+}$

$NH_4+(aq)+H_2O(aq)\rightleftarrows H_3O^{+}(aq)+NH_3(aq)$

Weaker the parent base = stronger the conjugate acid

Cations that are small, highly charged metals form weakly acidic solutions

• Cations such as $Al^{3+}$ and $Fe^{3+}$

• The hydrated form of the ion acts as a Brønsted-Lowry acid:

• The smaller and more highly charged the cation, the more acidic its behavior

The alkali and alkaline earth metals do not ionize water in this way

Salt Solution Behavior

Split the salt:

Salt→ Cation + Anion

Anion (conjugate base from acid)

Add H⁺ to recreate the acid:

• Strong acid → neutral anion (Cl⁻, NO₃⁻)

• Weak acid → basic anion (F⁻, CH₃COO⁻)

Cation (conjugate acid from base)

1. Try removing H⁺:

   • If it gives a weak base → cation is acidic

   • Example: NH₄⁺ → NH₃ → acidic

2. If no H or removal not possible:

   • Group 1/Group 2 metals / Ca²⁺ Sr²⁺ Ba²⁺ → neutral

   • Other small highly charged metals → acidic (Al³⁺, Fe³⁺)

Polyprotic Acids

• Definition: Acids that can donate more than one proton (eg, $H_2SO_4,H_3PO_4).$

• First dissociation ($K_{a1}$ ) is typically stronger than subsequent dissociations ($K_{a2}$ , etc.) because the first proton separates from a neutral molecule, while the second proton separates from an anion that holds the proton more tightly

Either

• The second proton contributes a negligible amount of H3O+ and $K_{a2}$ is ignored

• The second proton is significant and therefore you use the amount of original acid for the second reaction’s acid and resulting hydronium

Example: Sulfuric Acid

1. First dissociation (strong acid):

$H_2SO_4\left(aq\right)+H_2O\left(l\right)\to H_3O_{}^{+}\left(aq\right)+HSO_4\left(aq\right)$

• Assume complete dissociation for the first reaction

• If the initial concentration is 0.0100 M, then at the end of this step:

  • [HSO₄⁻] = 0.0100 M

  • [H₃O⁺] = 0.0100 M

Acid Strength and Molecular Structure

• The strength of binary acids depends on bond strengths and the polarity of the bond between hydrogen and the bonded atom.

• Strong vs Weak Acid Example HF is weak - HF is polar and the positive charge is on the proton, which makes it easier for the release of $H^{+}$

The strength of the bond affects the strength of the corresponding acid

• Stronger bond = Weaker acid (harder for the H+ to be released)

• Stronger electronegativity = Stronger acid

  • $HA\rightarrow H^{+}+A^{-}$ . If atom A is very electronegative, it pulls electron density away from hydrogen, making the hydrogen more positively polarized, easier for $H^{+}$ to be released

Oxyacids

Oxyacids contain a H atom bonded to an O atom. The O atom is bonded to another atom Y

The more electronegative Y $\rightarrow$ the more it weakens and polarizes the H-O bond $\rightarrow$ the more acidic the oxyacid

The more O atoms bonded to Y $\rightarrow$ the stronger the acid

• $H_2SO_4$ is stronger than $H_2SO_3$

• $HNO_3$ is stronger than $HNO_2$

Lewis Acids and Bases

Lewis Acids: Electron pair acceptors

Lewis Bases: Electron pair donors

$H^{+}\left(aq\right)+NH_3\left(aq\right)\rightarrow NH_4^{+}\left(aq\right)$

• $NH_3$ accepts a proton $\rightarrow$ donates an electron pair – acting as a base

• $H^{+}$ donates a proton $\rightarrow$ accepts an electron pair – acting as an acid

Molecules that can act as Lewis Acids/Bases include

• Molecules with incomplete octets that can accept an electron pair

• Molecules that can rearrange their electrons to create empty orbitals to be able to accept an electron pair

• Some cations can act as Lewis acids because of their empty orbitals