Acids and Bases Chapter 19

Ions in Solution  Aqueous solutions contain H + ions and OH - ions  If a solution has more H + ions than OH - ions it is acidic  If a solution has more OH - ions than H + ions it is basic  If a solution has the same number of OH - ions and H + ions it is neutral

Properties of Acids  Properties acids –  Produce H + ions when dissolved in water  taste sour  Turn Blue Litmus paper Pink  React with metals to produce Hydrogen Gas  good at dissolving things (food in stomach, teeth to form cavities, mineral deposits in coffeemaker)  Have a pH of 0 to < 7  Conduct electricity

Some Common Acids H 2 CO 3 Carbonic Acid Carbonated Water HC 2 H 3 O 2 Acetic Acid Vinegar H 3 C 6 H 5 O 7 Citric AcidFruits

Properties of Bases  Properties bases –  taste bitter  feel slippery  tend to produce OH - ions when placed in water  turn Red Litmus Paper Blue  Have a pH of 7 to 14  Conduct electricity

Some Common Bases NaOHsodium hydroxidelye KOHpotassium hydroxideliquid soap Ba(OH) 2 barium hydroxidestabilizer for plastics Mg(OH) 2 magnesium hydroxide“MOM” Milk of magnesia Al(OH) 3 aluminum hydroxideMaalox (antacid) Al(OH) 3 aluminum hydroxideMaalox (antacid)

Two models of Acids and Bases  Arrhenius  Bronsted-Lowry  Two similar but different explanations, both are correct

Arrhenius  Acid - produces H + in aqueous solution  HCl (g)  H + (aq) + Cl - (aq)  Base – produces OH - in aqueous solution  NaOH (s)  Na + (aq) + OH - (aq)  Explains most Acids and Bases, but not all

Acid-Base Neutralization Reaction  When an acid and a base are mixed together: acid + base   A salt and water are formed:  a salt + water  remember a salt is any ionic compound

Examples  HCl (aq) + NaOH (aq)  NaCl (aq) + H 2 O (l)  Acid+ BaseA Salt + Water  2HCl(aq) + Ca(OH) 2 (aq)  CaCl 2 (aq) + 2 H 2 O(l)  Acid+ BaseA Salt + Water

Bronsted-Lowry Model  Ammonia (NH 3 ) is a base.  It does not contain OH - in it.  So according in Arrhenius it isn’t a base.  But when placed in water it produces OH, so it must be a base.

Bronsted-Lowry Model  The Bronsted-Lowry model focuses on H + ion  Acid – H + donor  When placed in water Acids give H + away  HCl (g) + H 2 O (l)  Cl - (aq) + H 3 O + (aq)  H 3 O + (aq) is called the Hydronium ion  Base – H + acceptor  NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH -

Acid-Base Pairs  In Bronsted-Lowry Acids and Bases you must always have a pair  Every Acid must have a Conjugate Base  Every Base must have an Conjugate Acid HCl (g) + H 2 O (l)  Cl - (aq) + H 3 O + (aq) Acid Base Con Base Con Acid  HCl is an Acid and produces Cl -  H 2 O must be a Base and produces H 3 O +

More on Pairs  HI(s) + H 2 O(l)  H 3 O + (aq) + I - (aq)  AcidBase  C.A. C.B.  H 2 SO 4 (s) + H 2 O(l)  H 3 O + (aq) + HSO 4 - (aq)  Acid Base  C.A. C.B.

Strong Acids  Strong acid - reacts completely with water to produce ions; no molecules are left  Example: HCl + H 2 O  H 3 O + + Cl -  The Six Strong Acids: 1.HCl (hydrochloric acid) 2.HBr (hydrobromic acid) 3.HI (hydroiodic acid) 4.H 2 SO 4 (sulfuric acid) 5.HNO 3 (nitric acid) 6.HClO 4 (perchloric acid)

Strong acids:  HCl + H 2 O  H 3 O + + Cl- Remember, strong acids ionize completely in water. The reaction goes all the way to the right. A single arrow is used. There are virtually no HCl molecules left intact.

Strong acids:  HCl + H 2 O  H 3 O + + Cl- Looking at the equation above, with the single arrow, is Cl - a strong base or a weak base?

Weak Acids  weak acid – reacts only slightly with water to produce ions; mostly molecules left  HF + H 2 O  H 3 O + + F -  H 2 CO 3 + H 2 O  H 3 O + + HCO 3 -  Notice: Only one H comes off at a time

Strong Bases  Strong Base - reacts completely with water to produce ions; no molecules are left  Strong Bases:  Group 1 metals and Sr, Ba, and Ca with OH present  Examples NaOH, KOH, CsOH, Ba(OH) 2

Weak Bases  weak Bases – reacts only slightly with water to produce ions; mostly molecules left  Examples:  NH 3 + H 2 O   NH OH -  Fe(OH) 3 + H 2 O   Fe OH -

General rule: The conjugate base of a strong acid is a weak base. Similarly, the conjugate acid of a strong base is a weak acid.

One more reminder: “Weak” does not mean the same thing as “diluted.” HCl, for example is always a strong acid. If you add 1000 liters of water to it, it will be diluted, but still strong because what little there is will be completely dissociated.

Concentrated Vs Dilute  Concentrated means that there is a lot of Acid/Base molecules for the amount of water or that the Molarity is High  Concentrated HCl is 16 M  Dilute means that there are not many Acid/Base molecules or that the Molarity is Low  Diluted HCl would be 0.1 M

Naming Acids Look at the names of these acids – can you come up with the rule? H 2 SO 4 : sulfuric acid HNO 3 : nitric acid H 3 PO 4 : phosphoric acid

Naming acids Rule #1: If the acid comes from a polyatomic ion that ends in “ate,” the acid is named ____-ic. H 2 SO 4 : sulfuric acid (from sulfate) HNO 3 : nitric acid (from nitrate) H 3 PO 4 : phosphoric acid (from phosphate)

Naming Acids Rule #2: If the acid does not have oxygen in it, then name it… hydro + second element + ic Example: HCl is hydrochloric acid. What would HBr be? H 2 S?

Acid Nomenclature Flowchart

Acid Nomenclature Review No Oxygen  w/Oxygen An easy way to remember which goes with which… “In the cafeteria, you ATE something ICky”

HBr (aq) H 2 CO 3 H 2 SO 3  hydrobromic acid  carbonic acid  sulfurous acid Acid Nomenclature Review

Water as an Acid and a Base  amphoteric – describes substance that can act as an acid or as a base  Example: H 2 O (see previous Bronsted Lowry examples)  Arrhenius:  H 2 O  H + + OH -  Bronsted Lowry:  H 2 O + H 2 O  H 3 O + + OH -

Strength of Acids/Bases  pH Scale relates strengths of acids and bases  pH 0 to 7 – Acid  pH = 7 – Neutral  pH 7 to 14 – Base  pH can only be from 0 to 14  pH= -log[H + ]

BasicAcidicNeutral [H + ] pH Basic [OH - ] pOH

pH  pH= -log[H + ]  Used because [H + ] is usually very small  As pH decreases, [H + ] increases exponentially  Sig figs only the digits after the decimal place of a pH are significant  [H + ] = 1.0 x pH= sig figs

How to find pH  Punch in calculator: [H + ] or number, log, +/–  Example:  Find pH if [H + ] = 1.0 x M  pH = - log(1.0 x ) = 5.00  Punch in + / –, log, 1.0 EXP - 5, Enter

[H + ] Concentration  If you know the pH, you can determine the Molarity of H + Ions in the solution.  Since pH= -log [H + ]  Then [H + ] = 10 -pH  What is the Hydrogen Ion concentration of a solution that has a pH of 7.00?  [H + ] = = 1.0 x M

How to find [H + ]  Given pH, find [H]; Punch in pH or number, +/–, 10 x or 2nd log  Example: If pH is 9.0, find [H + ]  pH = -log [H + ]  Punch in 10 x or 2nd log, +/–, 9.0  [H] = 1.0 x M

What if it has OH - but not H + ?  14 = pH + pOH  1.0 x = [H + ] x [OH - ]  What is the pH of a solution that has a pOH of 5.00?  14 = pH or 14 – 5.00 = 9.00  What would the [H + ] be?  [H + ] = = 1.0 x M

pOH  Not all substances make [H + ], some make OH -  pOH= -log[OH - ]  14 = pH + pOH  Sig figs only the digits after the decimal place of a pH are significant  [OH - ] = 1.0 x pOH= sig figs  pH = 14 – 8.00 = 6.00

Finding OH- if you have H+ 1.0 x = [H+][OH-] 1.0 x = [H+][OH-] If you hydrogen ion concentration is 2.50 x 10 -5, what is the Hydroxide ion concentration? If you hydrogen ion concentration is 2.50 x 10 -5, what is the Hydroxide ion concentration? 1.0 x = [2.50 x ] x [OH-] 1.0 x = [2.50 x ] x [OH-] 1.0 x = [OH-] 1.0 x = [OH-] 2.50 x x = 4.0 x = 4.0 x

Indicators A solution or compound that changes color based on pH. Common Indicators: Universal Indicator (UI) Methyl Red Bromothymol Blue Methyl Orange