Electrolytes

Substances that, when dissolved in water,
produce aqueous solutions that will
conduct electricity

Strong electrolytes release many ions

Many ionic compounds

Weak electrolytes release few ions
Solutions

homogeneous mixtures in which one
substance is dissolved into another

the “solute” dissolves in the “solvent”

example: Kool-Aid -water is the solvent,
the drink mix is the solute
Molarity

•Molarity of a solution is equal to the number
of moles of solute divided by the number of
liters of solution
 M = mol/L

New symbol [square brackets]
 [H+] = molarity of H+ ions

The larger the molarity, the more concentrated
the solution
Concentrated

In a concentrated solution, the amount
of solute is large compared to the
amount of solvent it is dissolved in

Ex: juices, detergents
Dilute

In a dilute solution, there is much more
solvent than solute

Solutions are “diluted” by adding more
solvent

0.5M HCl is more dilute than 2.0M HCl

Acids and bases, as we use them in the lab, are
usually aqueous solutions.

Ex: when we talk about “hydrochloric acid”, it is actually
hydrogen chloride gas dissolved in water
 HCl(aq)

Concentrated acids have large molarities
 Ex: “conc.” HCl = 12M; H2SO4= 18M

In the lab, we usually use dilute solutions
 Ex: 1.0M HCl, or 0.1M H2SO4
What is an acid?

Many definitions are used

Arrhenius acid: a substance that
produces H+ ions in water

Then, H2O + H+  H3O+

H3O+= hydronium ion
Properties of Acids

React with most metals to produce H2(g)

react with carbonates to produce CO2

taste sour

damage living tissues

pH 0 -7

neutralize bases
Common Acids

Acid formulas –start with H

HCl –hydrochloric acid

H2SO4–sulfuric acid

HNO3–nitric acid

H3PO4–phosphoric acid

HC2H3O2–acetic acid

Also written CH3COOH
What is a base?

Commonly called “antacids”

•Arrhenius base:

•a substance that produces OH(hydroxide) ions when added to water
Common Bases

There are three common varieties of bases:

1)Hydroxide compounds (OH-)

ex: NaOH, Ba(OH)2

2)Carbonates (CO32-) and bicarbonates (HCO3-)

ex: Na2CO3, NaHCO3, CaCO3

3)Ammonia (NH3) and amines
Properties of bases

React with fats and oils to produce soap

feel slippery

taste bitter

damage living tissues

pH 7 -14

neutralize acids
Hydroxide bases

Release hydroxide ions directly into the
water

NaOH(s)  Na+(aq)+OH-(aq)
Carbonates and bicarbonates

React with water to produce hydroxide
ions

CO32-+ H2O  HCO3-+ OH-

HCO3-+ H2O  H2CO3+OH-
Ammonia and amines

React with water to produce hydroxide
ions

NH3+ H2O  NH4+ + OH-
Chemical Indicators
-chemicals that change colors when in an acid or base solution
Phenolphthalein
Litmus

acids = colorless

acids = red

bases = pink

bases = blue
pH scale

Used to indicate the acidity or “basicity” of
a solution

tells how strongly acidic a solution is…NOT
how strong an acid is!

Think pH as “parts H+”

•the lower the pH, the more H+’s, the more
acidic the solution
pH scale

The pH is the measurement of how
many H+’s are in the water…NOT a
measure of if the H+’s came from a
“strong” or “weak” acid!!!
pH scale

0-2
7=neutral

7.01 -10
 strongly acidic
 weakly basic

2 -4

10 -12
 moderately acidic
 moderately basic

4 –6.99

12 -14
 weakly acidic
 strongly basic
pH calculations
pH = -log
+
[H ]
=
+
[H ]
-pH
10
pOH calculation
pOH = -log
[OH ]
[OH-] = 10-pOH
pH and pOH relationship

In pure water at 25°C:
[H+] = 1x10-7M
 [OH-] = 1x10-7M
 Therefore, [H+] x [OH-] = 1x10-14
 And pH + pOH = 14

“THE BIG 5”
pH = -log [H+]
 [H+] = 10-pH
 pOH = -log [OH-]
 [OH-] = 10-pOH
 pH + pOH = 14

Autoionization of water

Water molecules can react with each other

H2O + H2O  H3O+ + OH-

At 25ºC, [H3O+] = [OH-]

[H2O] is a constant

Kw= [H3O+] [OH-] = 1x10-14

Let’s use [H+] instead of [H3O+]

Pure water is neutral
[H+] =[OH-] = 1x10-7M
 If [H+] > [OH-],the solution is acidic
 If [H+] < [OH-],the solution is basic

Arrhenius Neutralization

Works for the reaction of a strong acid
with a strong base

Remember –acid (or base strength) has to
do with how much of the acid (or base)
ionizes in water, not directly how many H+
or OH- are produced

100% ionization = “strong”
What is an acid?

Many definitions are used

Arrhenius acid: a substance that
produces H+ ions in water

Then, H2O + H+ H3O+

H3O+ = hydronium ion
What is a base?

Commonly called “antacids”

Arrhenius base:
 a substance that produces OH- (hydroxide)
ions when added to water
Arrhenius Neutralization
Hydroxide base- general form

Acid + Base  Salt + H2O

what’s actually happening?
 H+ + OH- H2O

Salt = the anion from the acid ……+ the
cation from the base
Arrhenius Neutralization

Examples with a hydroxide base
NaOH+ HCl  NaCl + H2O
 3H2SO4+ 2Al(OH)3Al2(SO4)3+ 6H2O

Arrhenius Neutralization
carbonate base- general form

Acid + Base  Salt + H2O + CO2

what’s actually happening?

2H++ CO32-H2CO3

H2CO3 H2O + CO2

Salt = the anion from the acid . . …+ the cation
from the base
Arrhenius Neutralization

Examples with a carbonate base
CaCO3 + 2 HCl  CaCl2 + H2O + CO2
 H2SO4 + NaHCO3  Na2SO4 + H2O +
CO2

Acid Strength

Compare the difference in these two
statements:

1)The more H+ ions in the water, the
more acidic the solution

2)The more H+ ions a compound
produces, the stronger the acid
Acid Strength

Strong acids release all of their H+ ions

[strong acid] = [H+]

Strong acids are strong electrolytes

Weak acids hold on to most of their H+ ions

[weak acid]>>>[H+]

Weak acids are weak electrolytes

Weak acids reach equilibrium with “neutralization” products
Don’t get confused!

A solution of a strong acid can be less
acidic that a solution of a weak acid!

IF: the strong acid solution is very dilute
and the weak acid is concentrated!