Properties

Acids

• Usually tastes sour.

• Can sting or burn skin.

• Conduct electricity.

• React with metals.

• Can neutralize a base.

• Has a pH<7

• Turns blue litmus paper red.

Bases

• Usually tastes bitter.

• Feels slippery on the skin.

• Conduct electricity.

• Do NOT react with metals.

• Can neutralize an acid.

• Has a pH>7

• Turns red litmus paper blue.

Chemical Properties and Identification

• All water solutions contain Hydrogen (H+) ions and Hydroxide (OH-) ions.

• An acidic solution contains more H+ ions than OH- ions.

  • Examples: HCl, HI, HBr

• A basic solution contains more OH- ions than H+ ions.

  • Examples: NaOH, KOH, Ba(OH)2

• The Hydronium ion is H3O+ and is the same indicator of an acidic solution as the concentration of H+ ions.

• When dissolved in water…

  • Acids always produce H+ ions

  • Bases always produce OH- ions

• Another indicator is Phenolphthalein.

• Turns colorless in acids.

• Turns bright pink in bases.

pH Scale

The pH scale is logarithmic, so an acid with a pH of 2 is ten times as acidic as an acid with a pH of 3 and 100 times more acidic than an acid with a pH of 4.

The pH scale ranges from 0 to 14.

• A weak acid is close to 7.

• A weak base is also close to 7.

• A strong acid is close to 0.

• A strong base is close to 14.

pH stands for the power of Hydrogen (H+)

Brønsted-Lowry Theory

Definitions

The theory states that an acid is always a hydrogen ion donor and a base is a hydrogen ion acceptor.

• Acids donate.

• Bases accept.

A conjugate acid is a substance produced when a base accepts an H+ ion and can donate its H+ ion to reverse the reaction.

A conjugate base is the substance produced when an acid donates an H+ ion and can receive an H+ ion to reverse the reaction.

Examples

HF + H2O → H3O+ + F-

• HF is the acid as it donates its H+ ion to H2O to create H3O+, and H2O is the base as it accepts the H+ ion from HF.

• H3O+ is the conjugate acid as it can donate its H+ ion to F- to reverse the reaction and create HF and H2O, and F- is the conjugate base as it can accept the H+ ion from the H3O+.

NH3 + H2O → NH4+ + OH-

• H2O is the acid as it donates its H+ ion to NH3 to create NH4+, and NH3 is the base as it accepts the H+ ion from H2O.

• NH4+ is the conjugate acid as it can donate its H+ ion to OH- to reverse the reaction and create NH3 and H2O, and OH- is the conjugate base as it can accept the H+ ion from NH4+.

Acids to Conjugate Bases

• HCl → Cl-

• H3PO4 → H2PO4-

Bases to Conjugate Acids

• H2O → H3O+

Titration

Definitions

A titration is a procedure done to determine the concentration of a solution with a known volume using a solution of known concentration using a burette. Concentration is the same as molarity.

A standard solution is the solution of known concentration in titration.

The Equivalence Point is when the concentration of H+ ions from the acid is equal to the concentration of OH- ions from the base to neutralize the acid.

The End Point is when there is a greater concentration of OH- ions from the base in the solution also known as the point at which the solution is one drop past the equivalence point. The solution will be light pink because of the phenolphthalein indicator.

Neutralization is when the concentration of the acid equals the concentration of the base, H+ = OH-.

Procedure

1. Clean and rinse the burette thoroughly before beginning a titration.

2. Using a funnel, pour about 10 mL of the titrant (the base) into the burette and roll it until covering the sides, then open the stopcock and allow the rinse to run out of the burette tip.

3. Fill the burette with the same solution using a funnel, and then drain it to exactly 0.0 mL and record the initial reading.

4. Prepare the solution to titrate in an Erlenmeyer flask with the appropriate volume with 2-3 drops of indicator, swirling to mix.

5. Begin titrating slowly by opening the stopcock and adding titrant to the flask drop by drop while simultaneously swirling the flask, avoiding splashing by putting the tip of the burette in the flask.

6. Continue titrating until the solution is a very light rose pink and lasts at least 30 seconds, one drop of titrant can darken the solution significantly.

7. Record the final reading on the burette and subtract the final reading from the initial reading to determine the volume in mL of titrant used.

pH and pOH

pH+pOH=14.00

pH formula

The pH scale is measured from 0 to 14; 0 being the most acidic with the most H+ ions and 14 being the most basic with the least H+ ions.

pH= 14.00-pOH

pH = -log[H+]

pOH formula

The POH scale is measured from 14 to 0; 14 being the most acidic with the least OH- ions and 0 being the most basic with the most OH- ions.

pOH = 14.00-pH

pOH= -log[OH-]

Calculating Concentrations of OH- and H+

When a substance is in brackets, it is defined as the concentration or molarity of the substance.

• [H+] is the molarity/concentration of H+ ions

• [OH-] is the molarity/concentration of OH- ions

[H+]=10^-pH

[OH-]=10^-pOH

Using the Formulas

pH = -log[H+] is the same as [H+] = 10^-pH

1. Divide both sides of the equation to isolate the logarithm. -pH = log[H+]

2. Using the properties of logarithms, the base of the log is 10, the output of the equation is [H+], and the exponent on base 10 is -pH.

3. So, the equation equals [H+] = 10^-pH

pOH = -log[OH-] is the same as [OH-] = 10^-pOH by solving in the same way.