Transcript than - Mrs. Walden`s Science Site

Chapter 5 and 17
Acids and Bases Introduction
What will make an
acid/base?
• General Rule:
• 1. If the oxide is covalent and a strong
bond holds the oxygen – acidic solutions
are produced
• Ex. SO3 + H2O  H2SO4
• 2. If the oxide is ionic – the compound will
produce a basic solution in water.
• Ex. CaO + H2O  Ca(OH)2
Properties of Acids:
• Sour taste
• Change color of indicators
• Some react with metals to produce H2
gas
• Are neutralized by the reaction with a
base
• Some conduct electricity
2 factors that determine
the strength of an acid:
• 1. Binary acids - The strength of a bond– the
stronger the bond, the weaker the acid (harder
to dissociate)
• 2. Oxyacids - The polarity of the bond – the
more oxygens – the more polar the molecule –
stronger the acid
• The more + charge of the metal cation in a
coordination compound – the stronger the acid –
increased polarity
• The more electronegative metal in an oxyacid –
stronger the acid due to increased polarity
Common Uses for Acids:
• A. Sulfuric acid – most commonly used
– in making of metals, paper, paints
• Attracts water – dehydration agent
• B. Nitric – rarely used – very unstable
– has a suffocating odor, stains skin,
burns
• Used to make explosives, rubber,
plastics
• C. Phosphoric – used in fertilzers,
detergents, ceramics, diluted in pop
• D. Hydrochloric – digestion, cleaning
agent, acidity in pools
• E. Acetic – (glacial acetic acid –
concentrated - will freeze at 17C)
• Vinegar is 4-8% acetic acid
• Used in plastics and food supplements
Types of Acids:
• Monoprotic – have one acidic H+ - ex.
HCl
• Diprotic – have 2 acidic H+ - ex. H2SO4
• Triprotic – have 3 acidic H+ - ex. H3PO4
• Polyprotic – acids that can donate more
than 1 acidic H+
• Organic – have a carbon backbone –
usually very weak – have only 1 acidic
hydgrogen
• Hydrohalic – acidic proton is attached
to a halogen – Ex. HCl or HF
Properties of bases:
•
•
•
•
Bitter taste
Change colors of acid/base indicators
Feel slippery
Are neutralized by the reaction of an acid –
produce salt and water
• Electrolytes
• Neutralization – when a strong acid and base
react they neutralize each other to form a salt
(ionic compound) and water
3 ways to define an acid/base
• 1. Arrhenius concept – acids produce H+
in aqueous solutions and bases produce
OH• Only applies to acids in aq solutions and
bases that contain OH-
2. Bronsted-Lowry Model
•
•
•
•
Acid is a proton donor
Base is a proton acceptor
Hydronium ion – H3O+
Polyprotic acids only dissociate one
acid at a time.
General Bronsted Lowry
Reaction:
• HA(acid) + H2O (base)  H3O+ (Conjugate acid)
+ A- (conjugate base)
• Conjugate base – everything that remains of the
acid after the proton is lost – will have a neg.
charge
• Conjugate acid – formed when the proton is
transferred to the base – (will have a + charge)
• Conjugate acid/base pair – 2 substances that
are related due the accepting/donating of a
proton.
• HA and A- (acid and its conjugate base) and
H2O and H3O+ (the base and its conjugate acid)
• The stronger the acid; the weaker the
conjugate base.
• The stronger the base, the weaker the
conjugate acid.
• Amphoteric (amphiprotic) – can act like an
acid or a base Ex. Water
• Autoionization – transfer of a proton from
one molecule to another of the same
substance to produce an acid and a base
3. Lewis Concept
• Lewis acid – electron pair acceptor (does
not have to be H)
• Lewis base – electron pair donor (does
not have to be H)
• Will form 1 product – acid –base adduct
• Look for bases that are anions or neutral
molecules that have lone pairs
• Look for acids that are cation or neutral
molcules with empty valence orbitals such
as B and Be
Acid-Base Indicators
• Compounds whose color changes
when the pH changes
• These are weak acids/bases
• Will be their original color in acidic
solution and a different color in a basic
solution as the indicator dissociates
• Universal indicators – have several
different indicators mixed together – will
show different colors at different pHs –
fairly accurate
pH meters
• Used if the exact pH is needed –
measures the voltage between 2
electrodes placed in the solution
• The voltage changes as the H+
concentration changes
Titrations
• Used to determine the concentration of an
unknown acid/base by a known acid/base
• Equivalence point (Stoichiometric point)
when the concentrations of the unknown
acid/base and the known acid/base are equal
– determined with an indicator or pH meter
• Endpoint – point during a titration where an
indicator changes color
• A good indicator’s endpoint matches the
equivalence point of the titration
How to determine the
equivalence point range:
• 1. Strong Acid with a Strong Base – pH
will be 7.00 at this point – neutral
• 2. Weak acids with a strong base – pH
will be greater than 7
• C. Weak bases with a strong acid – pH
will be less than 7
• D. Weak acid with a weak base beyond the scope of this class
pH Curve (Titration Curve):
• Plot of the pH of the solution as a
function of the amount of titrant added.
• Can use millimol (mmol) per milliliter to
describe titrations since the quantities
are usually small and burets are in mL
• Molarity = mmol/mL
2 important facts about
titration curves:
• 1. It is the AMOUNT of the acid, not the
strength that determines the amount of base
needed to reach the equivalence point.
• 2. The pH value at the equivalence point IS
affected by the acid strength. The weaker
the acid, the greater the pH at the
equivalence point.
• Standard solution – the known solution
• Primary standard – the highly purified
solid used to check the concentration
of the known solution
Steps on how to determine
the concentration of an
unknown through titration:
• 1. Write the balanced neutralization
reaction.
• 2. Determine the moles of the known
acid/base
• 3. Determine the moles of unknown
used during the titration.
• 4. Determine the molarity of the
unknonwn.
Equilibrium Constant – Ka
and Kb
• Strong acid and bases – equilibrium lies far to
the right – completely dissociaties at equilibrium
• will make a weak conjugate base/acid – water is
the main proton acceptor if an acid or proton
donor if a base.
• Large Ka if an acid or Kb if a base. K>1
• Weak acid or base – equilibrium lies to the left –
will hardly dissociate at equilibrium.
• conjugate base or acid is very strong –
conjugate base is the main proton acceptor or
conjugate acid is the main proton donor.
• Small Ka if weak acid or Kb if a weak base
• **Stronger the acid – the weaker its
conjugate base is**
• There is a competition taking place for
the H+ between water and the
conjugate base.
• If water is stronger – equilibrium lies
far to the right.
• If the conjugate base is stronger –
equilibrium lies to the left.
Acid – Base Properties of
Salts
• Salt – ionic compound – will break into
ions when they dissociate in water
• Salts that have cations of strong base
(Na+) and anions of strong acids (Cl-)
have no effect on the H+ concentration –
therefore, they are neutral – pH = 7.00
Salts of weak acids
• The conjugate base of a weak acid has an
affinity for protons – therefore conjugate
base affects the pH.
• A basic solution is formed if the anion of
the salt is the conjugate base of a weak
acid.
• Anions from polyprotic acids can act as an
acid or a base.
Salts of Weak Bases
• An acidic solution will be formed if the
anion is NOT a base and the cation is the
conjugate acid of a weak base – usually
only ammonium and its derivitatives
• If a salt contains a charged metal – will
from a complex ion
• Ex. Al makes Al(H2O)6+3 – this is a
conjugate acid
• Basic if written as [Al(H2O)5(OH)]+2
If both ions of the salt are
from weak acids/bases
•
•
•
•
Just compare the K values
1. If Ka > Kb – acidic
2. If Ka < Kb – basic
3. If Ka = Kb – neutral
• Ka * Kb = Kw
• Works for a weak acid and its
conjugate base
• Ka – weak acid; Kb is the conjugate
base
• pKa = -logKa
Predicting the direction:
• The reaction will always move from
the stronger acid/base to the weaker
acid/base
• If a weak acid and a weak base –
must compare the Ka and Kb values
of the conjugate acid and base.