Properties of Bases, Definitions of Acids and Bases, and pH Scale

Properties of Bases

  • Bases are proton \[H^+\] acceptors.

  • They have a pH greater than 7.

  • Bases taste bitter. An example is soap.

  • They affect indicators.

  • Bases feel slippery, like bleach, soap, or ammonia.

  • Bases neutralize acids.

  • They can be seen as hydroxide ($\OH^-$) producers in solution.

  • Hydroxide (\OH^-) can combine with hydrogen (H^+$) to form water (H_2O).</p></li></ul><h3collapsed="false"seolevelmigrated="true">HouseholdChemicalsandpH</h3><ul><li><p>VarioushouseholdchemicalsandtheirpHlevels;milkofmagnesiaismentionedasanantacidthatneutralizesstomachacid(hydrochloricacid)toproducesaltandwater.</p></li></ul><h3collapsed="false"seolevelmigrated="true">DefinitionsofAcidsandBases</h3><ul><li><p>Therearethreedefinitionsofacidsandbases:Arrhenius,Bronsted−Lowry,andLewis.</p></li><li><p>ThefocuswillbeonArrheniusandBronsted−LowryinChem1.</p></li></ul><h4collapsed="false"seolevelmigrated="true">ArrheniusDefinition</h4><ul><li><p>Anacidisasubstancethatproduceshydrogenions().</p></li></ul><h3 collapsed="false" seolevelmigrated="true">Household Chemicals and pH</h3><ul><li><p>Various household chemicals and their pH levels; milk of magnesia is mentioned as an antacid that neutralizes stomach acid (hydrochloric acid) to produce salt and water.</p></li></ul><h3 collapsed="false" seolevelmigrated="true">Definitions of Acids and Bases</h3><ul><li><p>There are three definitions of acids and bases: Arrhenius, Bronsted-Lowry, and Lewis.</p></li><li><p>The focus will be on Arrhenius and Bronsted-Lowry in Chem 1.</p></li></ul><h4 collapsed="false" seolevelmigrated="true">Arrhenius Definition</h4><ul><li><p>An acid is a substance that produces hydrogen ions (H^+$) in solution.

  • Arrhenius studied electrolytes and their effect on the conductivity of solutions.

  • Acids increase conductivity due to the presence of hydrogen ions.

Bronsted-Lowry Definition

  • An acid is a proton donor.

  • A base is a proton acceptor.

Lewis Definition

  • Deals with bonding and electron pairs centered around Lewis Structures.

  • A Lewis base brings the bond because it has a lone pair of electrons waiting for a hydrogen ion.

Bronsted-Lowry Acid-Base Definition and Conjugate Pairs

  • The Bronsted-Lowry definition views a hydrogen ion as a baseball that is thrown by the acid (proton donor) and caught by the base (proton acceptor).

Example

  • Ammonia (NH<em>3NH<em>3) + Water (H</em>2OH</em>2O) ⇌ Ammonium (NH4+NH_4^+) + Hydroxide (OH−OH^-)

  • In the forward reaction:

    • Water acts as an acid by donating a hydrogen ion.

    • Ammonia acts as a base by accepting a hydrogen ion.

  • In the reverse reaction:

    • Ammonium is the acid.

    • Hydroxide is the base.

  • Ammonia and ammonium are conjugates of each other, differing by a hydrogen ion.

  • The conjugates are on the product side of the reaction.

  • Ammonium is the conjugate acid, and hydroxide is the conjugate base.

Analyzing Conjugate Acid-Base Pairs

  1. Link conjugates: Find the same formula on both sides of the reaction that differs by one hydrogen and link them with brackets.

  2. Identify acid and base: Determine which reactant is the acid and which one is the base.

  3. Show hydrogen transfer: Draw an arrow showing the transfer of the hydrogen from the acid to the base on the reactant side.

Example Problems

  1. HBr + Water ⇌ Hydronium + Br

    • HBr and Br are conjugates.

    • Water and Hydronium are conjugates.

    • HBr is the acid, Br is the conjugate base.

    • Water is the base, Hydronium is the conjugate acid.

  2. General rules

    • Conjugate acid-base pairs differ by a hydrogen ion.

    • Bronsted-Lowry definition of acids uses the idea of conjugates.

  3. Organic derivatives of ammonia as weak bases are made when hydrogen is taken off and replaced with a CH3CH_3 group.

    • Trimethylamine is a derivative of ammonia (NH3NH_3).

    • Water in hydrolysis reaction acts like an acid if it donates a hydrogen ion or as a base by accepting a hydrogen ion.

Hydrolysis Reactions Involving Water

  • Hydrolysis reactions involve water as a reactant. Predicting products.

Example

  • HIO<em>3HIO<em>3 + Water ⇌ Hydronium + IO</em>3−IO</em>3^-

  • NH4+NH_4^+ + Water ⇌ Ammonia + Hydronium

Carboxylic Acids

  • Organic acids that contain the COOH group where the hydrogen on the end gets transferred during reactions.

  • R-COOH (R representing the rest of the molecule).

  • The COOH part = carboxylic acid part

  • The acidic hydrogen is on the end of the oxygen of the carboxylate group

    • The other hydrogens are non acidic

Hydrolysis with Water Acting as a Base

  • ClO<em>4−ClO<em>4^- + Water ⇌ HClO</em>4HClO</em>4 + OH−OH^-

  • Water acts like an acid.

  • ClO4−ClO_4^- acts like as base because it makes water turn into hydroxide.

  • Water can act as either an acid or a base. this is called amphoteric.

Acid-Base Properties of Water

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

  • Water undergoes self-ionization, producing hydronium and hydroxide ions.

  • This reaction occurs in all water samples, regardless of purity or temperature.

  • Therefore, water is not just H2OH_2O; it also contains hydronium and hydroxide ions in equilibrium.

pH Scale

  • pH stands for the power of the hydrogen ion concentration.

  • pH is numerically equal to the negative log of the concentration of the hydrogen ion.

  • pH=−log[H+]pH = -log[H^+]

  • Brackets indicate concentration, with units of molarity (M), or moles per liter.

  • The pH scale typically ranges from 1 to 14.

  • pH is defined as the negative log of the hydrogen ion concentration.

  • You can also determine the pOH, which is the negative log of the hydroxide ion concentration

  • pH + pOH has to equal 14.

  • A solution is neutral when the hydrogen ion concentration equals hydroxide concentration at 25 degrees Celsius.

  • If hydrogen ion concentration = 1×10−71 \times 10^{-7}, then pH = 7

  • You can only shortcut and estimate that the pH is equal to the exponent when the mantissa is 1.

Calculating pH

  • Use the formula to find the actual number

  • PH is equal to the negative log of 2×10−32 \times 10^{-3}

  • Pay attention to the amount of sig figs.

Significant Figures in pH

  • The number of significant figures in the concentration determines the number of decimal places in the pH value.

  • Example: If the concentration has one significant figure, the pH should be rounded to one decimal place.

Calculating pH of a Base

  1. Find the pOH using the formula: pOH=−log[OH−]pOH = -log[OH^-]

  2. Subtract the pOH from 14 to find the pH, using the formula pH+pOH=14pH + pOH = 14

Important Note

  • Neutrality depends on equal concentrations of hydrogen and hydroxide ions instead of pH=7pH=7. This is only true when the water is at room temperature. If water temperature is higher the amounts of hydronium and hydroxide is affected and thus changes neutral.

  • Hydrogen ion concentration is always equal to the concentration of hydronium ion because you need to account for hydrolysis.

  • Public service announcement: your body does a great job at maintaining pH!

Changes in pH

  • Changes in pH is logarithmic system so small changes are a big dang deal.

  • Every change from 7 to 6 of pH is a tenfold increase in the concentration of hydrogen ion.

Example

  • If Orange juice had a pH of 4