Transcript Acids and Bases - COHS IB and CP Chemistry

Acids and Bases
Properties of Acids
• They are corrosive
• pH range is 0 to 7, lower the pH, more
acidic the acid
• Taste sour
• Turns litmus paper red
• Often reacts with metals to produce H2
gas
• Often produces H+ in water. Example; HCl
 H+ + Cl-
Symbol used for acids
Properties of Bases
• pH range is from 7 to 14, higher the pH
more alkaline the base
• They are caustic
• Taste bitter
• Turns litmus paper blue
• Feels slippery
• Group 1 and 2 metals make common
bases. Example:
Na + H2O  NaOH + 1/2H2
The pH of some common acid and
bases
Common Acids
• Binary or Halide Acids
– HCl, hydrochloric acid: is stomach and pool
acid
– HF, hydrofluoric acid: used to etch glass
Common acids continued
• Oxyacids: acids with oxygen (more
oxygen- stronger the acid)
– HNO3, Nitric acid, used for making explosives,
when it reacts with metals it produces NO2gas
instead of H2
– H2SO4, sulfuric acid: used in car batteries,
used in many industrial process as a
dehydrator.
Chemical burn from H2SO4
• AHHHHHH SATAN!!!!!!! =]
Reaction of Cu with HNO3
Oxyacids continued
– H2CO3, carbonic acid, found in carbonated
sodas, and in rainwater. It is responsible for
the sour taste, or sting in sodas. It is also
found in blood to a very small extent.
– H3PO4, phosphoric acid, used as lime scale
remover, and as flavoring in the colorless
sodas like Sprite.
Acids continued
• Organic acids
– Carbon based acids.
O
– Structure ends in COOH
C-OH
– COOH is called a carboxyl group.
– CH3COOH, ethanoic acid or vinegar. Used as
a preservative and flavoring
Bases
• Group 1 bases
– Strongest bases
– NaOH, sodium hydroxide, and KOH, potassium
hydroxide, both of these bases are used to make
commercial cleansers and drain cleaners.
• Group 2 bases
– Weak bases,
– Ca(OH)2 calcium hydroxide, make limewater which is
used to test for CO2,,used to make antacids.
Bases
• Amines.
– This group has a -NH2 at the end of a
molecule
– These are formed from decomposing proteins
– They have a “fishy” smell
– NH3, ammonia, this is used as glass cleaner
and to make fertilizers.
– Examples CH3NH2, methyl amine,
CH3CH2NH2 ethyl amine. Etc.
Bases
• Carbonates (they are antacids)
– Have CO3 -2 in them
– This group produces CO2 when they react
with acids.
– NaHCO3 sodium bicarbonate, or baking soda
– CaCO3, calcium carbonate, or baking powder.
– MgCO3 and Al2(CO3)3 are used as antacids,
for example:
MgCO3 + 2HCl MgCl2(aq) + H2O(l) + CO2(g)
Arrhenius definition of acids and
bases
• Acids produce a hydrogen ion(H+) or
hydronium ion (H3O+) when they dissolve
in water. (An H+ ion is considered too
reactive to exist so an H3O+ ion is used)
• Acids dissolving in water producing an H+
– HCl  H+ + Cl– HNO3  H+ + NO3– H2SO4  2H+ + SO4-2
– CH3COOH  CH3COO- + H+
Carbonates are in antacids
Arrhenius definition continued
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Acids dissolving in water form a hydronium ion:
HCl + H2O  H3O+ + ClHNO3 + H2O  H3O+ + NO3H2SO4 + 2H2O  2H3O+ + SO4-2
CH3COOH + H2O  CH3COO- + H3O+
Note: H2O is written in the reaction when H3O+
is used, but not when H+ is used.
Arrhenius definition of bases
• Bases dissolve in water to produce an OHcalled a hydroxide ion.
– NaOH  Na+ + OH– Ca(OH)2  Ca+2 + 2OH– NH3 + H2O  NH4+ + OH– Note: H2O is written in the reaction only with
amines
Bronsted-Lowry def. of Acids/Bases
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Acids are proton donors
Bases are proton acceptors
Ex: HNO2 + H2O H3O+ +NO2Ex: NH3+ HCO3-   NH4+ + CO3-2
Ex: HCO3- + HSO4-   SO4-2 + H2CO3
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Determining the conj. Acid of a base
--Add a H+
Base
Conj. Acid
HCO3H2CO3
H2O
H 3O +
OHH 2O
• Determining the conjugate base of an
acid.
• --remove an H+
• Acid
Conj. Base
• H2O
OH• HSO4SO4-2
• NH4+
NH3
Continued Bronsted-Lowry
• Stronger the acid/ base, the weaker its
conjugate base/ acid.
• Amphoteric or Amphiprotic: A substance
which can be either an acid or a base.
• Example is H2O.
• Water as an acid: H2O + Cl-  OH- + HCl
• Water as a base: H2O + HCl  H3O+ + Cl-
Conjugate acid-base pairs
Definition of a Lewis acid and base
• Acids are electron pair acceptors in a dative
covalent bond.
• Bases are electron pair donors in a dative
covalent bond.
• These acids are reserved for those
molecules/ions which make dative covalent
bonds but are NOT ALREADY A BRONSTEDLOWRY ACID.
• Examples: NH3 is a base because the nitrogen
has a pair of electrons which can be used in a
dative bond. AlF3 can be a Lewis acid since Al
has a vacant pair of orbitals that can accept e-.
• Lewis acids will include substances that
are not considered typical acids.
• Examples:
– Some metal ions. Ex Group 3, Trans. metals
Strong acids and bases completely
dissociate in water
Strong and weak acids and bases
• Strong acid/base when it dissolves in H2O, nearly 100%
dissociates into ions.
• Ex: HCl H+ + Cl- nearly 100% of the molecules break up into ions
• Ex: NaOH Na+ + OH- nearly 100% of the molecules break up
(Remember: the H+ and OH- gives the properties associated with
Acids and Bases.)
▪ For weak acid and bases, when they dissolve in water only a
small % will dissociate to form ions.
CH3COOH(aq) CH3COO- (aq)+ H+(aq) Very little of the acid forms
ions. Same condition for weak bases.
▪ Both the strength and the concentration of the acid or base
determines its pH and its harmfulness. So when working with acids
and bases one needs to be concerned with
1. Is the acid or base strong or weak?
2. Is the acid or base concentrated or dilute?
WWCND
• How burgers are made
The pH equation
• The pH equation is: pH = -log [H+]
– Where [H+] is the molar concentration.
The pH scale
• Development of the pH scale
– Based upon the dissociation of water:
H2O  H+ + OH– Based upon the equilibrium of water:
Kw = [H+] [OH-]
• at equilibrium the concentrations of H+ and OH- are
each 1x10-7M (this was experimentally determined)
• Substitute these concentrations into the Kw
expression: Kw = [1x10-7][1x10-7]; this = 1x10-14
Development of the pH scale cont.
• According to the rules of equilibrium changing conc
does not change the Kw constant. So if acid is
added to water, the H+ conc goes up, and the OHgoes down, but K equals 1x10-14. The same thing
happens if a base is added to water. Knowing that
Kw is always 1x10-14, then
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–
–
The largest conc for either an H+ or OH – is 1M
The smallest conc H+ or OH – is 1x10-14 that we work with
the –log of 1M = 0 the low end of the pH scale
the –log of 1x10-14 = 14 the high end of the pH scale
Also, pure H2O has a conc of H+ = 1x10-7, the –log = 7,
pure H2O has a pH of 7 or it is neutral
Acid and Base Neutralizations
• General equation:
acid + base  salt+ H2O + energy
• Ex: HCl+ NaOH  H2O + NaCl
• CH3COOH + NaOH HOH +
NaCH3COO
• H2 SO4 + 2NaOH  2 HOH + Na2SO4
• HNO3 + NaOH  NaNO3 + H2O
• 2HCl+ Ca(OH)2  CaCl2 + 2H2O
Titration Curves for strong acid and
bases
Titration curve for strong acid and
bases
Explanation of a titration curve
• A titration involves an acid base reaction:
• Example NaOH + HCl  NaCl + NaOH
• All titration curves have the same shape but they do not
all start at the same pH
• Why this shape?
– Start with 1M HCl titrated with 1M NaOH
– Initial pH=O, titrate 10% of 1MHCl, 90% of HCl remains. [H+]=
.9M or pH= .05
– Titrate 50% of 1M HCl, 50% remains [H+]= .5M or pH= .3 (not a
great change in pH yet!)
– Titrate 90% of 1MHCl , 10% remains, [H+]=.1M or pH=1
– Titrate 100% of HCl, acid and bases are now equal, pH=7 the
cure shoots up until excess NaOH is added. Now curve
gradually increases again with the addition of more NaOH
Tiration curves with weak acids and
bases.
IB optional material: Equilibrium
and weak acids and bases
• Since weak acids and bases disassociate
partially:
– They are reversible and have a measurable
equilibrium.
– Their pH cannot be based upon the initial
molarity of the acid or base. An equilibrium
equation must be used to measure it.
The Ka or Kb (equilibrium constant)
of weak acids or bases