Transcript Acids and Bases

Acids and Bases
Types of Ions
• H+ ions – Hydrogen Ions
• OH- ions – Hydroxide Ions
Common Acids
• Hydrochloric acid (HCl)- stomach acid
• Carbonic acid (H2CO3) – carbonated drinks
•Nitric acid (HNO3)- explosives
•Sulfuric Acid (H2SO4) – used in industry
•Acetic acid (HC2H3O2)- vinegar
Properties of Acids
• Taste sour
• Produces gas when placed on a metal
or metal carbonate
• Turns blue litmus paper red
• Have more H+ hydrogen than OHhydroxide
Naming Acids
Acids all start with an H!!!
• Ending of name
as an ionic
compound ends
in:
-ide
-ite
-ate
• New name
becomes
–Hydro-root-ic
acid
–Root-ous acid
–Root-ic acid
Common Bases
• Sodium Hydroxide (NaOH) – soap
• Magnesium Hydroxide Mg(OH)2 –
antacid
• Group 1A or 2A cations w/ OH• Ammonia (NH3)- household
cleaners
Properties of Bases
• Feels slippery
• Tastes bitter
• Corrosive
• Burns skin
• Turns red litmus paper blue
• Higher concentration of OHhydroxide than hydrogen H+
Naming Bases
• Name the metal followed by hydroxide
regardless of how many hydroxides are there.
Practice Naming/Writing
Formulas
• Mg(OH)2
• H2SO4
• Hydrosulfuric acid
• Calcium hydroxide
• Aluminum hydroxide
• HBr
• LiOH
• Phosphoric acid
• Sulfurous acid
Strong Acids:
• HCl
• HBr
• HI
• H2SO4
• HNO3
• HClO4
• HClO3
Strong Bases:
• Lithium hydroxide
• Sodium hydroxide
• Potassium hydroxide
• Rubidium hydroxide
• Cesium hydroxide
• Calcium hydroxide
• Strontium hydroxide
• Barium hydroxide
Strong vs. Weak Acids and Bases
• Strong Acid/Base – An acid/base that
breaks apart completely when it is placed
in water. Are excellent conductors of
electricity
• Weak Acid/Base – partially ionizes and
therefore are poor conductors
Arrhenius Theory
• Acids contain hydrogen
• Bases contain hydroxide
• Neutralization:
• Acid + Base  Water + Salt
• Hacid + BaseOH  HOH + BaseAcid
• Hydrogen from the acid bonds with hydroxide from base
combine to form water. The other 2 ions bond to form a
salt (ionic bonds)
HNO3 + NaOH  H2O + NaNO3
According to Arrhenius Identify
Acid or Base
• H2S
• NaOH
• Al(OH)3
• H2CO3
The Amphoterism of Water
• Amphoteric: The ability to act as an ACID or
BASE depending on what it is around.
• While other substances are able to do this as
well, WATER is the most common
Self Ionization
• When two water molecules combine
to form one OH- and one Hydronium
atom H3O+
Bronsted Lowry Theory
Sheep Theory
• An acid is a proton donor
• A base is a proton acceptor
• Conjugate Acid – Base Pairs
– A conjugate acid of the base → Base + H+
– A conjugate base of the acid → Acid − H+
According to Bronsted Lowry
Determine Acid, Base, Conjugate
Acid and Conjugate Base
Kw: Water Constant
• Value of the equilibrium constant expression
for the self ionization of water
• Kw = [H+] x [OH-]
• If it is water H+ and OH- will be equal so if H =
1.0 x 10-7 then OH = 1.0 x 10-7 making water
1.0 x 10-14
• Kw = 1.0 x 10-14
Kw: Water Constant
• If H+ is larger then it is an acid
• If OH- is larger then it is a base
H+
OH-
4.7 x 10-5
7.0 x 10-7
2.3 x 10-12
1.0 x 10-7
A/B
What is a log?
A logarithm (log) is the exponent to which a
power of 10 must be raised to get the number.
-log 1 x 10-3
log = 3
pH
• The negative logarithm of the hydrogen ion
concentration
pH = -log[H+]
• pH can not be less than 0 nor greater than 14
• If it is below 7 it is an acid, above 7 a base, and
7 is neutral
pH
7
0
Acids
Bases
14
Ion Concentrations from pH
• Antilog (-pH) = [H+]
pOH
• Negative logarithm of the hydroxide ion
concentration
pOH = -log[OH-]
• pOH less than 7 are basic, greater than 7 are
acidic, and 7 is neutral
pH
7
0
Acids
Bases
Bases
Acids
14
14
7
pOH
0
Ion Concentration from pOH
• Antilog (-pOH) = [OH-]
• pH + pOH = 14
H+
OH-
pH
pOH
5.4 x 10-9
2.7 x 10-3
3.3
1.6
A/B
Titration
• Using a KNOWN concentration to figure
out an UNKNOWN concentration of an
acid or base
1) What is the concentration of a phosphoric
acid solution if 25.0 mL are exactly neutralized
by 20.0 mL of 2.000 M KOH solution?
KOH + H3PO4  K3PO4 + H2O
2) How many mL of a 0.1 M calcium hydroxide
solution are needed to completely neutralize
10.00 mL of 0.250 M HCl solution?
Ca(OH)2 + HCl  CaCl2 + H2O
RICE diagram
• RICE diagrams are needed when asked to find
pH or pOH
• Reaction
• Initial
• Change
• Excess
3) In the titration of 55 ml of 0.5 M HCl with 0.6 M
NaOH, what is the pH after 50 ml of the NaOH
solution has been added?
HCl + NaOH  H2O + NaCl
4) In the titration of 18 ml of 0.3 M HCl with
0.25 M NaOH, what is the pH after 20 ml of the
standard NaOH has been added?
HCl + NaOH 
H2O +
NaCl
5) In the titration of 30 ml of .75 M HCl with 0.5
M NaOH, what is the pH after 35 ml of NaOH
solution has been added?
HCl + NaOH 
H2O +
NaCl
Vocabulary to Add
• Amphoteric: substance that can act as an acid
or base (Water most common example)
• Saponification: mixing a base and fat to make
soap
• Indicator: any substance that gives a visible
sign, usually by a color change, of the
presence or absence of a threshold
concentration of an acid or base solution
Electrochemistry
OIL RIG
Oxidation is loss (of e-)
Reduction is gain (of e-)
LEO the lion goes GER
• Lose electrons = Oxidation
• Gain electrons = Reduction
Balancing Acid Redox Reactions
• Split the reaction into two halves
• Balance the common element on the two
sides.
• Balance the Oxygen by adding water to the
opposing side
• Balance the hydrogen by adding H+ ions to the
opposing side
• Balance the charge by adding electrons
Redox Reactions in Acid
• Al + MnO4-1  Al3+ + Mn2+
MnO4-1+ Fe2+  Fe3+ + Mn2+
Cr2O7-2 + HNO2  Cr+3 + NO3-1
C + H+ + SO42-  CO2 + SO2
Balancing Basic Redox Reactions
• Split into two half reactions
• Balance the common ion
• Balance the charge of the common ion by
using electrons
• Balance the overall charge by using OH• Balance the hydrogen by adding H2O
Basic Redox
• Calculate the charge of P in PO4-3
• Calculate the charge of Cr in Cr2O7-2
Redox Reactions in Basic Solutions
• Mg + OCl-  Mg(OH)2 + Cl-
• Cl- + Cr2O7-2  Cr3+ + ClO2-
PbO2 + SeO32-  PbO + SeO42-
MnO4- + NH3  NO3- + MnO2