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Binary acids - Milford High School

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Binary acids - Milford High School Powered By Docstoc
					Acids bases & salts
            Objectives
State the Bronsted-Lowry definition of
acids and bases
Identify the common physical and
chemical properties of acids and bases
Explain what dissociation constants
indicate about an acid or base
Use experimental data to calculate a
dissociation constant
  Properties of acids/bases
Taste – acid comes from latin meaning
sour or tart/bases are bitter – soap
Touch most dilute acids feel like water
although they sting on broken skin bases
feel smooth, soothing and slippery except
in your eyes (soap)
Reactions with metal acids react bases
do not react
      Properties continued
Electrical conductivity water is poor –
HCl is good, NaOH is good both are
electrolytes
Indicators turn color – acid turns litmus
paper from blue to red base turns from red
to blue
Neutralization reaction between an acid
and a base get salt and water (double
replacement)
       Arrhenius definition
An acid is a substance that dissociates in
water to produce hydrogen ions.
A base is a substance that dissociates in
water to produce hydroxide ions
A salt is an ionic compound formed from
any cation other that H+ and any anion
other than OH- or O-2
      Arrhenius continued
Acids react with metals to produce H2 gas
Mg +2H+ -> Mg+2 + H2
oxidation reduction reaction
Bronsted – Lowry definitions
An acid is any substance that can donate
H+ ions
A base is any substance that can accept
H+ ions a Bronsted-Lowry acid is a proton
donor and a base is a proton acceptor
          Hydronium Ion
H+ is strongly attracted to the electrons of
surrounding water molecules
H+ + H2O -> H3O+
More correct HCl + H2O -> H3O++ Cl-
In this case HCl is the Bronsted-Lowry
acid and water is the base
We still describe a solution of HCl as
acidic!
 NH3 + H2O -> NH4+ + OH-
Ammonia is the H  + acceptor

water is the H+ donor (acid)


  Amphoteric a substance that
  can act as either an acid or a
              base
  Conjugate Acid-Base Pairs
NH3 + H2O  NH4+ + OH-
In 1 direction water is the acid in the reverse
reaction it is the base.
These cmpds become conjugate acids and
conjugate bases when HCl loses an H+ ion to
become its conjugate base Cl- when the
conjugate base of water is the hydroxide ion OH-
When ammonia gains H+ to become its
conjugate acid NH4+ and the conjugate acid of
OH- is H2O
         Conjugate pairs
A pair of compounds that differ by only
one H+ ion such as H2O and OH- or NH3
and NH4+ are called conjugate acid base
pairs
NH3 + H2O  NH4 +      + OH-
Base acid conj. Acid conj. base
Determining the strengths of acids
           and bases
A strong acid HCl readily transfers
hydrogen ions to water to form H3O+
If you place 1M of HCl in 1 liter of water
you would form 1 M H3O+ ions and 1 M Cl-
ions
              weak acids
A weak acid does not readily transfer H+
ions
1 mole of acetic acid in 1 liter of water only
.4% of the acetic acid molecules would
form H3O+ and C2H3O2-. Which means
that 99.6% of the acetic acid molecules do
not dissociate.
To show a strong acid from a weak
        acid use arrows
 HCl + H2O  H3O+ + Cl-
 HC2H3O2 + H2O <-> H3O+ + C2H3O2-
   Strong and Weak Bases
The most widely used commercial base is
CaO. When CaO is dissolved in water the
O-2 ions react completely with H2O to form
OH- ions.
O-2 + H2O  2OH- use a single arrow
Strength of conjugate acid – base
              pairs
The stronger the acid the weaker its
conjugate base.
The stronger the base the weaker its
conjugate acid.
The acid dissociation constant Ka
Weak acid HA
HA + H2O <--> H3O+ + A-

Keq =
For a 1 M solution of a typical weak acid may be only
.007% of the water molecules react. Move the water to
the left side of the equation

The higher the Ka the more the reaction goes to the right
. The greater the Ka the stronger the acid
Weak acids have a Ka less than 1
           Diprotic acids
2 step dissociation
Base dissociation constant Kb
The base dissociation constant is a
measure of the strength of a base
Do calculations of dissociation constants
            Objectives
Explain what most acidic hydrogen atoms
have in common
Explain what most bases have in common
Describe how acids are named
Naming and identifying acids and
            bases
1st all H’s are not acidic CH4
As a rule an acidic hydrogen already has a
slight positive charge while is it part of a
molecule. (It is in the positive side of a
polar covalent bond)
Usually bonded with O, N, or a halogen
          3 types of acids
Binary acids - H and 1 other element usually
6A or 7A
Strong HCl, HBr, and HI
Weak HF, H2S and H2Se
Oxy acids contain H, O and 1 other element
H2SO4, HNO3 and H3PO4
Carboxylic Acid – organic acids COOH group
Acetic acid HC2H3O2 vinegar
                Bases
A Bronsted-Lowry base always contains
an unshared pair of electrons NH3 attracts
H+
Anions: remember conjugate bases HCl
Cl-
Weak Cl- Br- I- NO3- HSO4- CIO4-
Strong O-2 OH- PO4-3 and CO3-2
Amines N has an unshared pr. Of
electrons
   Naming Acids and Bases
If the name of an anion ends in ide the
name of the acid that produces it includes
the name of the anion, a hydo prefix
Hydrochloric acid – all binary acids
If the name of an anion ends in ate use
and ic ending Nitric acid, carboxylic acid
Anion ends in ite SO3-2 sulfite ion H2SO3
sulfurous acid

				
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posted:11/30/2012
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