Acids and Bases

Acids  Taste sour  Begin with H  Found in many foods and drinks  Turn blue litmus paper red  pH  Corrosive  Forms H + (or H )ions in solutions HydrogenHydronium

Bases  Bitter  End in OH  Turn red litmus paper blue  pH  Found in many cleaning products  Slippery  Corrosive  Forms OH - ions in solution Hydroxide

Acids and Bases  Neutral: H + = OH -  Acidic: H + > OH -  Basic: H + < OH -  ↑H + = ↓OH - = more acidic = ↓ pH ac idic  ↓H + = ↑OH - = more basic = ↑ pH

Water = Neutral H 2 O = HOH HOH → H + + OH - Free Hydrogen ion bonds with water molecule to form Hydronium ion H + and H 3 O + used interchangeably

Hydronium Ion = Hydrogen Ion  Self ionization - two water molecules react to form a hydronium ion (H 3 O + ) and hydroxide ion. H 2 0 → H + + OH -

pH scale  Shows the strength of acid or base on a scale of  Numbers below 7 = acids…the lower the number, the more acidic  Numbers above 7 – bases…the higher the number the more basic.

pH and pOH  pH + pOH = 14  If the pH = 2, what is pOH  If the pH = 4, what is the pH  If the pOH = 7, what is the pH

Calculating pH  Formula: pH = -log [ H+ ]  You can calculate pH by finding the negative logarithm of the concentration of hydrogen ions.

Calculating pH  A solution contains 1.0 x mol/L of H + ions, what is the pH of this solution?  Formula: pH= -log [H + ]  pH = -log (1.0 x )  pH = 8 = Base

Calculating pH  A solution contains 3.5 x M of H + ions, what is the pH of the solution?  Formula: pH= -log [H + ]  pH= -log ( 3.5 x )  pH = 4.5 = acid

pH and Water  Water is amphoteric ; it can act as both an acid and a base in an aqueous solution.  Water contains an equal number of H + and OH - ions. H2OH2OH+H+ + OH -

Ion Product Constant of Water  K w is the ion product constant for water.  Represents the equilibrium for the self ionization of water.  Formula: K w = [H + ][ OH - ]

 [H + ] = 1.0 x  [ OH - ]= 1.0 x  What is the K w ?  K w = (1.0 x ) x (1.0 x ) K w = 1.0 x This is a constant on your STAAR Chart

 The H - concentration of an aqueous solution is 1.0 x M. What is the OH - ion concentration? K w = [H + ][OH - ] K w = 1.00x List the Knowns H+ = 1.0 x x = 1.0 x [OH - ] plug into formula [OH - ] = 1.00 x / 1.0 x = 1.0 x mol/L M

 The OH - concentration of an unknown solution is 2.4 x What is the H + concentration of the solution? Is the unknown solution acidic, basic, or neutral?  K w = [H + ][OH - ]  1.0 x = [H + ][ 2.4 x ]  [H + ] = 1.0 x / 2.4 x  [H + ] = 4.16x  -log(4.16 x )= pH= 10.4  = Basic solution

Acid Base Reactions  Acid + Base = neutralization reaction  Acid + Base → water + salt (always)  Salt = (+) ion from base & (-) ion from acid  Positive ions are always listed first HA + BOH → HOH + B + A -

Arrhenius  Swedish Chemist Svante Arrhenius created a model for acids and bases in 1883.

Arrhenius Model- Acid  HCl (g) H + (aq) + Cl - (aq) Acid is a substance that contains hydrogen and ionizes to produce hydrogen ions in aqueous solution.

Arrhenius Model- Base Base is a substance that contains a hydroxide group and dissociates to produce a hydroxide ion in aqueous solution. NaOH (s) Na + (aq) + OH - (aq)

Bronsted- Lowry Model  Danish chemist Bronsted and English chemist Lowry proposed a model that focuses on the Hydrogen Ion  An Acid is a hydrogen-ion donor  A Base is a hydrogen-ion acceptor

Ionization  The Bronsted-Lowry Model also shows if and acid or base is strong based on ionization.  Strong acid- completely ionized  Weak acid- partial ionization

Strength and Concentration  Strength – how completely it ionizes Strong – ionizes completely or almost completely Weak – ionizes partly  Concentration Concentrated - a lot of acid/base in water. Dilute – a little acid/base in water.

 12 M HCl is a strong acid with a high concentration  Adding 6L of water to this solution would do what to the solution: strong acid, dilute solution  Vinegar has acetic acid, which is weak, in low concentration = dilute  12 M acetic acid would still be weak, because it only partially ionizes, but it would be a concentrated solution, because there is a lot of acid dissolved in a little water.

Strong Acids  Since strong acids are completely ionized they produce the maximum number of ions.  Strong acids are good conductors  Reaction only moves in one direction, represented with an arrow in one direction.

 HClH + + Cl -  HBr H + + Br –  H 2 SO 4 H + + HSO 4 _

Weak Acids  An acid that ionizes partially in dilute aqueous solutions  Produce fewer ions, so they are poor conductors  Reactions move both directions until equilibrium is reached, represented by an arrow in both directions

 HFH + +F -  H 2 S H + +HS -  H 2 CO 3 H + + HCO 3 -

Conjugate Acid  The species produced when a base accepts a hydrogen ion to form an acid Conjugate Base The species that results when an acid donates a hydrogen ion to form a base.

Conjugate acid – base pairs  2 compounds with the same chemical formula, but the acid of the pair will have 1 more H NH 3 & NH 4 - H 2 SO 4 & HSO 4 - H 2 O & H 3 O +

Bronsted-Lowry Model  NH 3 + H 2 O → NH OH -

Precipitate Reactions  When two compounds come together to form an aqueous compound and a solid compound.  2NaOH(aq)+CuCl2(aq)  2NaCl( aq )+Cu(OH) 2 ( s )  KI(aq) + AgNO 3 (aq)  KNO 3 ( aq ) + AgI( s )  Use your STAAR chart to check solubility  If insoluble – compound will precipitate or settle out of solution as a solid

Oxidation-Reduction Reaction  A reaction in which electrons are transferred from one atom to another  2KBr(aq) + Cl 2 (aq)  2KCl(aq) + Br 2 (aq)  The chlorine on the left steals electrons from the bromine in KBr to become KCl and Br 2 on the right.

Oxidation- Reductions Reaction

Remember Acid-Base Reaction  Form SALT + WATER  Mg(OH) 2 + 2HCl  MgCl 2 + H 2 0 base + acid  salt + water  Salt = any ionic compound made up of a cation (+) from a base and an anion (-) from an acid

Identify the following reaction: as 1) precipitation, 2) oxidation-reduction, or 3) acid-base  2K + Br 2  2KBr  H 3 N + CsOH  Cs 3 N + H 2 O  MgCl 2 + Li 2 CO 3  MgCO 3 + LiCl Look on STAAR chart to see if either compound Is insoluble

Titration  Use known solution ( standard solution) to find the concentration of an unknown solution  Drop by drop process  Endpoint – point of color change of indicator  When neutralized

Buffers  Resist changes or swings in pH  Blood pH approx 7.4  Fatal if fall or rise more than 0.3 pH units  Buffers in your blood prevent big changes when, for example, you eat an orange (citric acid)