Download CHAPTER 2

ST 410/EST 404
THE MATERIAL WORLD
CHAPTER 2 NOTES
CHAPTER 2
Molecules
(pp. 40-50)
1 What is a molecule?
Only a few elements, such as _________________ and __________________, exist in their pure
form on Earth. Most atoms combine with atoms of other _____________________ to form
___________________________.
 A molecule is a group of __________ or more atoms that are _____________________
bonded together.
Examples: -
O2
-
_____________
-
H2O
-
_____________
-
NaCl
-
_____________
Why do atoms tend to bond with other atoms?
______________________________________________________________________________
______________________________________________________________________________
 Noble gases (Group VIII) have a _____________ valence shell; therefore they are extremely
_________________ and rarely ________________ with other elements.
 Halogens (Group VII) have _______ valence electron, so they need to ____________ one
electron to acquire the electron configuration of the nearest noble gas.
 Alkali metals (Group I) have only ________ valence electron, so they all tend to _________
that electron to resemble a noble gas.
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Table 2.4: THE TENDENCY OF GROUP A ELEMENTS TO GAIN OR LOSE ELECTRONS
Group #
IA
IIA
IIIA
IVA
VA
VIA
VIIA
VIIIA
Element
example
Li
Be
B
C
N
O
F
Ne
# valence
electrons
Tendency
Octet Rule: The tendency of elements to acquire the configuration of the noble gas
__________________ to them in order to have ____________ electrons in their valence shell.
(Exceptions: Li, ______, and ______ acquire the configuration of ______________ and thus
follow the ___________ rule.)
*Special case: ___________________ - depending on the circumstances it can _________ its
only electron or it can __________ a second electron.
1.1 IONS
In general atoms are electrically ____________________ (equal # of _____ & _____ )
An ion is an atom that has become electrically ___________________ by ________________ or
_____________________ one or more electrons.
ION FORMATION IN METALS
 Since alkali metals (Group I) have only ________ valence electron, they all tend to
_________ that electron when forming ions. When this happens, they acquire a charge of
_________.
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http://www.nios.ac.in/images/5.1.gif
 Since alkaline earth metals (Group II) have ________ valence electrons, they all tend to
_________ _____ electrons when forming ions. When this happens, they acquire a charge of
_________.
All metals ____________ their valence electrons when forming ions and thus form
_____________________ charged ions (CATIONS).
ION FORMATION IN NON-METALS
 Since halogens (Group VII) have ________ valence electrons, they all tend to _____________
_________ electron when forming ions. When this happens, they acquire a charge of
_________.
Because non-metals all have ______ or more valence electrons, they all _____________
electrons when forming ions and thus form _____________________ charged ions (ANIONS).
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1.2 The Nature of Chemical Bonds (EST ONLY)
Most atoms, except those of noble gases, have a natural tendency to ____________ or
_________ electrons in order to fill their outer shells. When two atoms come together, they will
either ___________________ or ___________________ their valence electrons to become
_________________.
 A __________________ _________________ is the union of two atoms through the
___________________ or _______________________ of one or more electrons.
There are _______________ main types of chemical bonds: ___________________ bonds and
______________________ bonds.
IONIC BONDS
 An IONIC BOND is usually the result of a transfer of one or more _________________
from one atom (usually a ___________) to another atom (usually a ________________).
The formation of an ionic bond represented with Lewis structures
 http://www.clickandlearn.org/Gr9_Sci/atoms/bonding.htm
In the Lewis dot diagram above we see that when the sodium atom comes in contact with a
chlorine atom, the sodium atom gives up an _____________________. Both atoms thus acquire
an electron configuration similar to that of a ______________ gas. The sodium atom becomes a
___________________ ion (Na+), and the chlorine atom, a negative ion (Cl-). Since positive and
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negative charges _______________ each other, the positive sodium and the negative chloride
ion come together to form a _______________________ compound.
In the space below, draw the Rutherford-Bohr Atomic Model showing the formation of an ionic
bond between Magnesium and Bromine (MgBr2).
COVALENT BONDS
Molecular oxygen (O2), ammonia (NH3), and methane (CH4) are examples of the type of
bonding where an electron _____________________ reacts with another ________________.
 A COVALENT BOND is the result of the _____________________ of one or more
electron _______________ between two ______________________ atoms.
When molecular fluorine (F2) is formed, each atom _____________ an electron with another
fluorine atom so they both have the electron configuration of __________, the nearest Noble gas.
In the Lewis structure, the shared electron pair is ___________________; in the ball and stick
model, it is represented by a _______________.
F
F
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Sometimes two atoms share than one ________________ pair. In molecular
____________________ (O2), each oxygen atom needs ____________ more electrons to
achieve the configuration of a ________________ gas, so ___________ oxygen atoms tend to
share two electron pairs. Oxygen atoms are linked in a _________________ bond. Triple
bonds, between atoms of other elements (such as _________________) are also possible.
Rutherford-Bohr Atomic Model for O2
Lewis Structure for O2
Ball and Stick Atomic Model for O2
Electron pairs are not always shared ____________________. Some atoms have a ___________
force of attraction for electron pairs than others. In a _______________ molecule, the
____________atom attracts the electrons more than the two _____________________ atoms do.
This causes a certain degree of ________________ polarity, with the oxygen atom carrying a
slightly _____________ charge, and the hydrogen atoms, a slightly __________________ one.
The charges are so small that they do not turn water into an _______________ compound. The
unequal covalent bonds are referred to as __________ ___________ ___________.

symbol for a
partial charge
O
H


H


Diatomic molecules
Some elements do not exist in nature as
individual atoms. Such atoms come in
pairs like socks and jeans. They are
diatomic molecules (made up of 2 atoms).
To recall which elements are diatomic,
just remember this simple phrase:
“I Have No Bright Or Clever Friends”
I2, H2, N2, Br2, O2, Cl2, F2
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1.3 The Rules of Chemical Notation and Nomenclature (EST ONLY)
Naming Binary Ionic Compounds
 A binary compound is a compound made up of _______ different elements. (bi = _______)
1. Name the metal first
2. Add the suffix –ide to the name of the non-metal
Examples:
NaCl __________________________________________
CaF2
__________________________________________
AgBr __________________________________________
Na3N __________________________________________
KI
__________________________________________
ZnO
__________________________________________
Mg3P2 __________________________________________
Writing Formulas for Binary Ionic Compounds
Use the CROSS-OVER RULE!!!!
The Cross-Over Rule involves writing the charge on each ion as a superscript and then crossing
the numbers over and writing them as subscripts without the + and -. Don’t worry, it’s super
easy!
 Remember to reduce to lowest terms!
Example: Write the molecular formula for magnesium bromide.
Mg2+
Br-
MgBr2
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Let’s do a few more:
Aluminum oxide
Sodium fluoride
Calcium sulfide
_____________
_____________
_____________
Stock System for Naming Ionic Compounds Containing Multi-Valent Ions
Some transition metals can form ions with two different charges. Because of the existence of two
different ions for these metals, we need a naming system that will enable us to distinguish one from the
other. The system we will use is the Stock System. The Stock System involves writing a roman numeral
after the name of the metal to distinguish it from its other ion.
Examples:
Fe2+ = __________________________________
Fe3+ = __________________________________
Pb2+ = __________________________________
Pb4+ = __________________________________
Cu+ = __________________________________
What do we do when we’re faced with naming this: CuCl2?
Is it copper (I) chloride or copper (II) chloride?
We have to do the Cross-Over Rule in reverse!
Cu? ClCuCl2
Since the charge on chlorine is ______ and there are _____ chloride ions in the formula, that makes 2 × -1
= _____, so the copper has to have a charge of _______ in order for the compound to be neutral.
So the name of this compound is ___________________________________________.
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Let’s try a few more:
Fe2O3
PbS
MnO2
___________________________
___________________________
___________________________
Writing Formulas for Ionic Compounds Containing Multi-Valent Ions
Use the CROSS-OVER RULE just like you do for a regular ionic compound!
Example: Write the molecular formula for chromium (III) chloride.
Cr3+
Cl-
CrCl3
Try these:
Nickel (II) bromide
___________
Gold (III) oxide
___________
Mercury (II) sulfide
___________
Naming Binary Covalent Compounds
 Use prefixes to indicate the number of atoms of each type.
 DO NOT use the prefix “mono” on the first element in the formula.
Table 2.18: Prefixes Indicating the Number of Atoms of an Element in a Binary Covalent Compound
Number of Atoms
Prefix
Number of Atoms
One
Six
Two
Seven
Three
Eight
Four
Nine
Five
Ten
Prefix
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Examples:
CO2
__________________________________
N2O4 __________________________________
CO
__________________________________
SF6
__________________________________
PCl3
__________________________________
P4O10 __________________________________
Exceptions to naming covalent compounds
 Some covalent compounds have common names and are not named according to the rule
above. (You must memorize the table below!)
Formula
Name
Formula
H2O
CH3OH
NH3
C2H5OH
CH4
C6H12O6
C3H8
C12H22O11
C4H10
H2O2
Name
 Hydrogen compounds (ex. HCl, H2S, etc.) DO NOT take prefixes!
HCl
____________________________
H2S
____________________________
POLYATOMIC IONS (EST ONLY) p. 44
 A ___________________________ ION is a group of _______ or more chemically
bonded atoms that has become electrically _________________________ by
_____________ or ____________________ one or more electrons.
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Table 2.10: Examples of Common Polyatomic Ions
Chemical Formula
Name
CH3COO-
Chemical Formula
Name
OHAmmonium
Nitrate
Bicarbonate
Nitrite
CO32-
PO43Chlorate
SO42-
CrO42-
Sulfite
NOTE: Your teacher will ask you to memorize some of the most common polyatomic ions.
Naming ionic compounds containing polyatomic ions (non-binary ionic compounds)
Examples:
CaCO3
________________________________
Mg3PO4
________________________________
NaOH
________________________________
Na2SO4
________________________________
HCN
________________________________
NaHCO3
________________________________
Writing formulas for ionic compounds containing polyatomic ions (non-binary ionic
compounds)
 Use the CROSS-OVER RULE!
 If there is more than one polyatomic ion in the formula, you must put brackets around it.
 Never change the subscripts of a polyatomic ion! Ex. Ca3(PO4)2 ≠ Ca3P2O8
Example: Write the chemical formula for magnesium hydroxide.
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Ammonium nitrate
Potassium chromate
Aluminum sulfate
___________
___________
___________
2 Properties of solutions
Sometimes atoms and molecules can combine without undergoing a ______________ reaction to
form a ______________. Since not _____________ bonds need to be _____________, the
different substances that make up a mixture can be ___________________ using physical
_______________ ________________.
 A solution is a ___________________ mixture (consisting of at least one ________ and
one ________________) whose component substances (solids, _______________ or
gases) cannot be _________________, even with the aid of a magnifying instrument.
 Solute: The component of the substance that is ______________ in the other.
Examples of solutes include salt, sugar, colouring and alcohol.
 Solvent: The substance in which the solute ________________. Examples of solvents
include water, alcohol and acetone.
 Aqueous Solution: A solution in which the solvent is _________________.
Water is the universal solvent because:

It dissolves many substances. Molecules with _______________ bonds and molecules
with a certain polarity dissolve easily in water. ________________ molecules, such as
oil, rarely dissolve well in water.

It is _______________ (pH = 7)

_________________________

Odourless

_________________________

Doesn’t react
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2.1 SOLUBILITY
 SOLUBILITY is the maximum amount of _________________ that can be
_________________ in a certain volume of ____________________.
Factors that affect solubility:

_______________ of the solute

_______________ of the solvent

_______________________ affects solubility of gaseous solutes

________________________ (solids tend to become more ____________ as solvent
temperature rises while gases tend to become less _______________ as solvent
temperature rises)
Figure: Solubility of Carbon Dioxide in Water as a Function of Temperature
HTTP://WWW.ENGINEERINGTOOLBOX.COM/GASES-SOLUBILITY-WATER-D_1148.HTML
 SEE Appendix 2 on p. 516 for a list of the solubility (and other characteristic properties) of many
common substances.
2.2 CONCENTRATION
 The CONCENTRATION of a solution is the_______________ of _______________
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in a given amount of _________________. It is the ratio of the quantity of solute to the
quantity of the solution.
DILUTION AND DISSOLUTION
The concentration of a solution can be varied in different ways.
Change
Effect on the concentration
Dilution (_____________ of solvent)
Dissolution (addition of ________________)
___________________ (reduction of solvent)
Expressing the Concentration of Aqueous Solutions
1. Concentration: Number of grams of solute per liter of solution (g/L)
Application: What mass of NaOH is needed to prepare 500mL of a 4 g/L NaOH solution?
2. Mass-Volume Percent: Number of grams of solute per100 mL of solution, expressed as a
percentage (% m/V)
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Application:
You have 24g of sugar to prepare a 6%m/V sugar solution. What volume of solution will you
make?
3. Volume Percent: Number of millilitres of solute per 100 mL of solution, expressed as a
percentage (% V/V)
Application:
a)
You have 50mL to prepare a 6%V/V alcohol solution. What volume of solution will
you make?
b) You add 75 mL of acetone to 1205mL of water. What is the concentration of the
solution in %V/V?
4. Mass Percent: Number of grams of solute per 100 mL of solution, expressed as a
percentage (%m/m)
Application: What is the mass of NaCl in 400g of an 8%m/m brine solution?
CONCENTRATION IN PPM
When the amount of solute in the solution is very small, the concentration can be expressed in
_____________ _________ _________________.
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 The CONCENTRATION in PPM (“ __________ ____
_______________”)
is the number of parts of solute in a _______________ parts of solution.
1 ppm =
1g
1mg
=
= 1 mg/L
1000000 g 1000 g
VERY IMPORTANT!!!!
 X%(m/V) = X g/100mL
 Example: 5%(m/V) = 5g/100mL = 5000 mg/0.1 L = 50 000 mg/L or 50 000 ppm
MOLAR CONCENTRATION (AKA Molarity) (EST ONLY)

Expressed as moles per liter (__________________)
MOLAR CONCENTRATION corresponds to the number of ____________ of dissolved
solute particles in a ___________ of solution.
 Molar concentration is symbolized by placing the _______________ formula for the
measured substance inside _________ brackets. Example [NaCl] = 0.5 mol/L means the
___________ concentration of sodium chloride solution equals _____________ mol/L.
Molar concentration of a solution can be calculated using the formula:
The units for molarity are mol/L or M (in this case, M is the unit for molar concentration.
MC is the symbol for molar concentration.)
Example: Suppose 58.5 g of sodium chloride are dissolved in 500 mL of solution. Calculate
the molar concentration of this solution following the method illustrated on page 54.
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More Practice
a) If 20g of KNO3 is dissolved in enough water to make 500mL of solution, what is the
molar concentration of the solution? (Answer 0.4M)
b) What mass of CaF2 is needed to prepare 250mL of a 0.1mol/L solution?
(Answer 1.95 g)
 The following is not covered in Chapter 2 of OBSERVATORY. It is, however,
important information that you need to know for your lab exam in June.
HOW TO PREPARE A SOLUTION
How would you prepare 250mL of a 20 g/L solution of cobalt (II) chloride, CoCl2?
1. Calculate the mass of solute needed.
2. Using an electronic balance weigh out the mass of the solute needed.
3. Pour a small amount of water into the volumetric flask.
4. Pour the solute into the flask.
5. Add water up to etched line. Use a pipette near the end (bottom of meniscus should be on
the etched line)
6. Cap the flask, invert the flask and shake. Repeat 3 times.
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DILUTING SOLUTIONS
To dilute a solution is simply to add ____________ to a more concentrated solution. Even after
you do this, the amount of _________________ is the same in both the original concentrated
solution and the new diluted solution.
Recall that C = m/V, so m = C  V
DILUTION FORMULA
mc = mass of solute in concentrated solution
mc = m d
md =mass of solute in diluted solution
Cc Vc = Cd Vd
Cc = concentration of the concentrated solution
(in %m/m, %V/V, %m/V, or g/L)
Cd = concentration of the diluted solution
(in %m/m, %V/V, %m/V, or g/L)
Vc = volume of concentrated solution
(in mL or L)
Vd = volume of diluted solution
(in mL or L)
Some tips to remember:
Cc  Cd (ex. 30%V/V  6%V/V)
Vd = Vc + Vwater added
Vc  Vd (ex. 5 mL  25 mL)
Vwater added = Vd - Vc
Example 1: You have a bottle of 5%m/V bleach solution. You need to prepare 500 mL of a
2% m/V bleach solution. What volume of concentrated bleach will you use?
Example 2: What is the final concentration when 100 mL of water is added to 200 mL of a
12 g/L salt solution?
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Example 3: What volume of water must be added to 500 mL of a 30%V/V hydrogen peroxide
solution to dilute it to 6%V/V?
Example 4: Using 300 mL of a 45 g/L solution of lithium chloride, LiCl, a student must prepare
a 15 g/L solution of lithium chloride. What is the volume of the resulting solution?
2.3 ELECTRICAL CONDUCTIVITY
Pure water does not __________________ electricity. How then does it does carry and electric
current? It is the substances _____________________ in the water that conducts the electricity.
 An __________________________ is a substance that when dissolved in water,
allows an _____________________ _________________ to flow through the solution.
 The ELECTRICAL ______________________ of a solution is a measure of its
ability to allow ______________________ ____________________ to flow through it.
 A ________________________________ is a substance that is soluble in water but
DOES NOT conduct electricity.
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ELECTROLYTIC DISSOCIATION

When an _____________________________ is dissolved in water, it separates into two
__________ of opposite charge, one ____________________ and one
_______________________.

This separation, known as _______________________ _________________________,
is a _____________________ change.
The chemical equation for the electrolytic dissociation of sodium chloride is:
 The H2O over the arrow indicates that the change takes place when the
_______________ is placed in water.

The ions formed during this process _______________________ electricity.
A non-electrolyte does not conduct electricity when dissolved in water because it does not
_________________________.
 All covalent compounds (except acids) are non-electrolytes.
Here’s what happens when you dissolve methanol (CH3OH) in water:
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THE STRENGTH OF ELECTROLYTES (EST ONLY)
To determine whether a solute is an electrolyte:
1. Dissolve it in ___________________
2. Place two _______________________ in the solution
3. Connect to a power supply and a light bulb
4. Does the bulb light up?
Strong Electrolytes
Substances that dissociate
Weak Electrolytes
Substances that only
Non-Electrolyte
Substances that ________________
___________________ (100%)
___________________ dissociate
produce ions when dissolved in
when dissolved in water.
when dissolved in water.
water.
CaCO (s) –>50% → Ca2+(aq) + CO32-(aq)
C12H22O11(s) –H2O→ C12H22O11(aq)
100 molecules
100 molecules
NaCl –H2O→ Na+(aq) +
100 molecules
Cl-(aq)
100 ions + 100 ions
50 ions + 50 ions
100 molecules
no light
 The strength of an _______________________________ is the degree to which it
dissociates into ions. The higher the degree of dissociation, the
____________________ the electrolyte.
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TYPES OF ELECTROLYTES
Acids

Bases

Found in fruit juices, ____________
Found in many ______________
____________, & gastric juices
products and in some

pH ______________ than 7
_______________________

Taste ________________
medication. Blood and _______ water

Turns blue litmus paper ____________
are also slightly basic.

Neutralize _________________

pH ________________ than 7

Release H+ ________ when dissolved

Taste _______________________
in water

Turn _____________ litmus paper blue
Molecular formula often begins with

Neutralize _________________
the symbol for a _________________

Feel _________________ to the touch
atom followed by a nonmetal (HCl,

Dissolve _______________ and oils
HNO3, HSO4, HF, H2CO3)

Release ________________ ions (OH-

 Acetic acid
(aq))

(___________________________) is
when dissolved in water
Molecular formulas begins with a metal
and usually ends in “OH”:
the exception to this rule.
Examples of acid solutions:
NaOH, KOH, NH4OH, Ca(OH)2,
1. Hydrochloric acid: HCl → H+(aq) + Cl-(aq)
2. Nitric acid: HNO3 → H+(aq) + NO3-(aq)
3. Sulfuric acid: H2SO4 → 2H+(aq) + SO42-(aq)
Al(OH)3
 Exceptions: Alcohols! CH3OH,
C2H5OH  these are NOT bases and
NH3 is a base yet it does not end in OH
Examples of basic solutions:
1.
2.
3.
Sodium hydroxide: NaOH → Na+(aq) + OH-(aq)
Potassium hydroxide: KOH → K+(aq) + OH-(aq)
Magnesium hydroxide: Mg(OH)2 → Mg2+(aq) +
2OH-(aq)
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Salts

Constitute a ____________ class of substances that figures ________________ in human
diets

Most are made up of a _______________ and one or more ________________________.

pH = _________ (___________________)

Have __________ effect on litmus paper
 A SALT is a substance produced by the ________________ bonding of a
____________________ ion and a ______________________ ion (other than H+ and
OH- ions).
Examples of salt solutions:
1.
2.
3.
4.
sodium chloride: NaCl,
potassium bromide: KBr
calcium chloride: CaCl2
silver nitrate: AgNO3
 Not all salts dissolve easily in water.
2.4 pH
Acidic, basic and neutral solutions can be distinguished by their _________. A solutions pH
can be measured using a pH _______________________ or a pH _______________.
THE pH SCALE

ranges from __________ to ____________

if the pH the solution is __________________

if the pH = 7, the solution is _______________________

if the pH the solution is ______________________

the pH scale is ____________________ , which means that a difference of one unit
between two substances actually indicates that one of the substances is _________ times
more acidic than the other.
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Examples:
a) What is the pH of an acetic acid solution that is 100 times more diluted than an acetic
acid solution with a pH of 2?
____________________________________________________________________
b) What is the pH of an ammonia solution that is 1000 times more concentrated than an
ammonia solution with a pH of 10?
___________________________________________________________________________
Figure 1: The pH of some common substances
http://islandwood.org/kids/stream_health/Data/pH_scale.jpg
State whether the following substances are acidic, basic, or neutral:
Susbtance
Acidic, Basic, Neutral?
Vinegar
Bleach
Coke
Detergent
Sea Water
END OF CHAPTER 2 NOTES FOR ST
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MORE ON pH (EST ONLY)
The pH of a solution is actually an indication of the concentration of __________________ ions
(_____) in that solution.
How is pH related to hydrogen ion concentration?
pH = -log [H+]
where [H+] = hydrogen ion concentration in mol/L
So, if the hydrogen ion concentration in a solution is 1 10-3 mol/L, then the pH = _______.
Likewise,
[H+] = 10-pH
So, if the pH of a solution is 8, then the hydrogen ion concentration is _____________________.
 See Table 2.30 on page 61 of your textbook.
THE pOH SCALE
The complete opposite of the pH scale, it communicates hydroxide ion concentrations, [OH-], in
a wide variety of substances.
pOH = -log [OH-]
where [OH-] = hydroxide ion concentration in mol/L
So if the hydroxide ion concentration in a solution is 1 10-6 mol/L, then the pOH = ________.
Likewise,
[OH-] = 10-pOH
So, if the pOH of a solution is 10, then the hydroxide ion concentration is __________________.
How are pH and pOH related?
pH + pOH = 14
So if the pH of a solution is 5, then the pOH is ______.
END OF CHAPTER 2 NOTES FOR EST
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