Laboratory Techniques for Environmental Technicians
Environmental Monitoring
& Technology Series
Chemical, Forensic, Food & Environmental Technology
Laboratory
Techniques
For Environmental Technicians
Trainee Learner Resource
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Laboratory Techniques for Environmental Technicians
Chapter 1: Introduction and Induction
Practicals and theory
Introductory Laboratory Practice - details and back-ground information
Workplace protocols
Chapter 2: The language and structure of Chemistry
What is Chemistry?
Practical: Observation Skills
Classification schemes
Practical work
Practical: Physical versus chemical change
Bohr model of the atom
The Periodic Table & Atomic Structure
Chemical Formulae
Types of compounds
Organic chemicals
The Mole Concept
Solutions
Calculating the molarity of a solution
Dilution of solutions
Practical: Metal colours in a flame
4
4
4
5
6
6
6
7
12
17
19
27
28
31
38
43
47
48
50
52
Chapter 3: Safety in the Laboratory
53
Safety in the Laboratory
Laboratory hazards
Laboratory rules and regulations
Practical 1.1 Laboratory Layout
Clean-up of spills
Practical 1.5 Safety in the laboratory
Dangerous goods 'Class' labels
53
54
56
57
58
59
62
Chapter 4: Basic Laboratory Equipment
Basic Laboratory Equipment
Chapter 5: Introduction to Material Handling
Material transfer
Practical work 5.1 Basic material handling techniques
Volume measurement techniques
Accurate and Approximate Volume Measurement
Pipetting Techniques
Burette Techniques
Volumetric Flask Techniques
Measurement Of Mass
Practical work 5.2A Introductory weighing task
Other Basic Laboratory Procedures
Practical work 5.5A Simple filtration
Practical: 4.4 Heating devices
64
64
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70
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Laboratory Techniques for Environmental Technicians
Practical:
Application of heating equipment
87
Chapter 6 The dreaded calculations
89
Laboratory Data and Error Analysis
Classification of Errors
89
90
Chapter 7: The why’s of sampling
Practical:
Practical:
Validation of sampling
Sampling Equipment
Chapter 8: Solution Preparation
Practical 5.4A Introductory Solution Preparation
Chapter 9: Basic environmental laboratory testing
Laboratory Measurements
Practical work 7.1 pH measurements
Measurement of density
Practical:
7.7 Density
Chapter 10: Gravimetric analysis
Gravimetric Analysis
Practical: Analysis of a mercury chloride and determination of its empirical formula.
Practical: Gravimetric determination of sulfate in bore water 8.6
Practical: Experimental Investigation of combustion of magnesium
Practical: Gravimetric determination of Ni by precipitation with dimethylglyoxime.
Chapter 11: Volumetric Analysis
Types of Standards used in Titration
Titration Calculations
Practical: Practice titration
Practical: Ethanoic acid content of Vinegar
Practical: Sodium Carbonate Content of Washing Soda
Practical: Chloride by titration with silver nitrate
Practical: Determination of hardness in water
Chapter 12: The Open Ended Problem
Open ended project
95
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107
109
110
112
112
112
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115
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119
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126
129
132
138
141
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154
154
Assessment 1
156
Assignment 1
Assignment 2
Assignment 3
Assignment 4
156
157
158
159
Assessment 2
160
Assignment 1
160
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Laboratory Techniques for Environmental Technicians
Chapter 1: Introduction and Induction
The Chemical, Forensics, Food & Environmental Technology (CFFET) section of Hunter TAFE
has developed this manual to help online students in the area of Laboratory Practice. This
manual is designed to be used in conjunction with the text “Practical Laboratory Skills”
(1995 Harcourt Aust) by Krajniak, Barker & Fullick. It is not a requirement to purchase this
text as the required information will be provided throughout the chapters that follow. The
CFFET team gratefully acknowledge the permission of the authors to reproduce text and
diagrams in these notes.
Laboratory Practice is a general term used by this teaching section to describe standard
operating procedures in a typical science laboratory. Activities in this manual have been
designed to help candidates learn and practice skills necessary to be deemed competent in
some of the competency standards described in the Laboratory Operations Training Package
– MSL09. Laboratory Practice covers all the specified activities, defined in TAFENSW
curriculum, to develop general science laboratory competencies.
Practicals and theory
This manual only provides brief introductory notes for each activity and it is expected that
the teacher and the textbook will provide more extensive details of basic theory and the
contexts in which the theory may be applied. The additional information will be provided
during the intensive session.
The practical manual provides result sheets to record and report the laboratory activities,
observations and measurements made which summarise the learner’s work. These
completed result sheets will become part of the evidence of the learner’s skills needed to
assess competency. There will be selected activities or responses required prior to the
intensive session these will be identified by a box requiring an answer that will be sent off to
the trainer.
Introductory Laboratory Practice - details and back-ground
information
Staff details: Head Teacher ………………….Dr. G. Fullick……………………………………
Class Teacher ………Adam Samuelson; Daniel Solomon; Denise Hatton
Technical staff……depends on availability……………………
Specific details relating to working in the Science Building of Hunter Institute will be
provided at the first session during the intensive session. The following page highlights the
workplace protocols that apply to the Edgeworth David Building, Hunter Institute. These will
be discussed in the induction first session
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Laboratory Techniques for Environmental Technicians
Workplace protocols
1. Safety Procedures
(a)
(b)
(c)
(d)
(e)
(f)
(g)
Consult Safety Data sheets* and method of
analysis for advice on hazards and
precautions to be taken
Wear appropriate PPE
Use fume hood etc as necessary
Maintain tidy workspace
Exercise care not to endanger other people
Observe emergency procedures
Report spillages and all accidents
4.
Testing
(a)
Refer to workplace procedures manual
for standard method
Conduct tests according to workplace
procedures
Clean up spills promptly
Record results according to workplace
procedures, without alteration
Calculate results, checking against
expected values and correcting errors
Trouble shoot basic problems with
procedure or equipment which have led
to atypical results
(b)
(c)
(d)
(e)
(f)
2. Recording and Reporting
(a)
(b)
(c)
(d)
(e)
(f)
(g)
(h)
(i)
(j)
(k)
(l)
(m)
(n)
(o)
Register samples into laboratory system
Label samples
Record which tests the sample should
undergo
Record sample description, compare with
specification, record and report
discrepancies
Record calibration results for
instruments/equipment in tables and/or
charts, following quality system
Keep records of calibration status and
calibration schedule for instruments /
equipment
Report faulty equipment
Keep records of solutions prepared, by
expected use-by date, and by name of
person who prepared them
Record results legibly, and chart when
required to identify trends
Interpret trends
Identify and report atypical results
promptly to appropriate personnel
Record approved results into workplace
system
Comply with quality system
Report all accidents and potential hazards
Maintain confidentiality of workplace
information
3.
Sample Handling
(a)
(b)
Maintain sample integrity
Prepare sample and standards for test
5.
Equipment and Reagents
(a)
(b)
(g)
Set up equipment and reagents
Check calibration status of equipment;
calibrate if necessary
Monitor shelf-life of working solutions
Prepare solutions when necessary, label
and log into laboratory register
Clean and care for test equipment and
work space
Dispose of faulty equipment or
quarantine it for repair
Store unused reagents
6.
Wastes
(a)
(b)
Minimise generation of wastes
Collect, sort and dispose of wastes in
accordance with procedures
7.
(a)
Environmentally specific
Equipment not in use is turned off at the
power switch
Lights are off when laboratory is
unattended
Fume cupboards are off when not in use
Non-conformances with environmental
standards are reported
(c)
(d)
(e)
(f)
(b)
(c)
(d)
* SDS’s were up until January 1 2012 referred
to as MSDS (Material Safety Data Sheets)
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Laboratory Techniques for Environmental Technicians
Chapter 2: The language and structure of Chemistry
What is Chemistry?
Scientists like to group the things they observe. It provides a way of organising and putting
their observations in order. I has evolved over many thousands of years using the
observation skills of many people. Both quantitative and qualitative information is
important. What is the difference between the two terms:
Quantitative information
This is any information that has measurement attached to it. It could be a temperature
measurement for example 25°C or a length for example 25 mm or the number of trees in a
100 m2 area is 11.
Qualitative information
This is information that has no exact amount associated. It is hot or it is cold or the frog is
about 10 cm.
Many important scientific discoveries have been the result of someone noticing something
that many may not have observed or treated as insignificant. It is important to develop both
your quantitative and qualitative observation skills.
Observation skills are particularly important in aspects of laboratory work. We use all
senses to assist in working safely and also in noting aspects of our practical work
Practical: Observation Skills
Take a candle and:
Make as many observations as you can about your candle.
Type you answer here
Leave blank for assessor feedback
Light your candle and record all your observations
Type you answer here
Leave blank for assessor feedback
Blow out your candle and record your observations
Type you answer here
Leave blank for assessor feedback
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Laboratory Techniques for Environmental Technicians
Classification schemes
List any classification systems for science with which you are familiar.
Type you answer here
Leave blank for assessor feedback
Matter is a good starting point for classification and hence the study of chemistry.
Matter, Change and Energy
Chemistry is a natural science that deals with the composition of matter and the changes it
undergoes. Which of the following are examples of matter?
◗ concrete
◗ acetone (propanone) vapour
◗ heat
◗ sound
◗ air
◗ light
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Laboratory Techniques for Environmental Technicians
Matter is anything that has mass and occupies space. Matter can exist as solid, liquid, gas
and plasma.
Some Classification Schemes
Solids/ Liquids/ Gas
Matter can exist in three states, solid, liquid or gas (and recently a fourth state, plasma has
been recognised).
Gas
Liquid
Solid
Describe how the particles are arranged in each of the three states identified in the diagram
above
Type you answer here
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Laboratory Techniques for Environmental Technicians
Characteristics of the states of matter
Gas
Liquid
Solid
Particle
position
widely
spaced, less closely spaced, closely
move freely
some collisions
vibrate
position
Shape
varies
varies-takes shape definite
of container
Volume
varies
definite
definite
Compress-
Easily compressed
slightly
Negligible
rapid
slow
negligible
in
packed,
fixed
ibility
Diffusion
In which state of matter do the following exist at room temperature and pressure?
◗ diamond
◗ oxygen
◗ cooking oil
◗ mercury
◗ carbon dioxide
◗ sugar
◗ Concrete
◗ Oxygen
◗ Air
◗ Nitrogen
Type you answer here
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Laboratory Techniques for Environmental Technicians
Pure/ Impure
A pure substance is either an element or a compound. A pure substance is identified by its
constant physical properties.
Elements are substances composed of only one type of atom. They have fixed physical
properties. The known elements have been arranged into a table known as the Periodic
Table.
Depending on the element the atoms could be arranged in one of the following;
◗ Monatomic, composed of only one atom for example helium
◗ Diatomic, composed of two atoms of the same element for example oxygen
◗ Triatomic, composed of three atoms of the same element for example ozone
◗ Polyatomic, composed of many atoms of the same element for example diamond,
graphite
Where elements can exist in more than one form they are referred to as allotropes of an
element for example:
◗ Carbon in the form of graphite, diamond, bucky balls
◗ Oxygen in the form of oxygen gas and ozone gas
◗ Phosphorus in red and white forms
List any elements that you are familiar with. Use a periodic table to find the position on the
table and also the symbol. (A periodic table is available in the Appendix file)
Element
Symbol
Oxygen
O
Answer
Answer
Answer
Answer
Answer
Answer
Element
Symbol
Answer
Answer
Answer
Answer
Answer
Answer
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Laboratory Techniques for Environmental Technicians
Use a periodic table to fill in the comical ‘Gold Dust Kid’ activity
An Elemental Tale: The Gold Dust Kid
The Kid mounted his trusty steed, old (B) ……………. His shooting (Fe)…………… strapped to his
side, he headed out for the bright (Ne)…………… lights of Sabattus, aiming to rob the
Litchfield stage. There was sure to be a load of precious
(U)……………… aboard, and
probably (K)……………………too. Inhaling a deep breath of (O) ………………. he coughed on the
(S)……………….. from the nearby hills. Since the (Hg) …..…………….was climbing, he quenched
his thirst with some H2O, tasting the (Cl) ……………… from (Ca) ……………….. deposits built up
over the years of riding the (Zn) ……………. trail. Overhead a (He) …………..-filled balloon
floated in the breeze; the sun beat down like burning (P)…………… Soon he spotted the stage,
guarded only by a sheriff with a (Sn)……………badge. “Halt”, he yelled, “or I’ll fill you full of
(Pb)……………. The sheriff drew his gun, but alas, was too slow. The Kid’s gun blazing like
flaming (Mg)…………..……… did the (Cu)……………….. in. Anyone who drew on the Kid should
know his life wasn’t worth a plugged (Ni)…………….. A (Pt) ………………… blonde riding beside
the (Al) …………………..-framed coach rode for her life when the Kid pulled out some (N)
………………… compounds, preparing to blow the safe to atoms.
Suddenly a shout rang out, “Hi Ho (Ag)…………….. and a masked man on a white horse raced
across the (Si) ……………….. sands like (Na) ……………….. skittering on water. A (H) …………………
bomb would not have stopped the lawman; the Kid had met his doom. The rest of his life
was to be spent behind (Co) ………………. steel bars, a warning to all who flirt with danger.
Your first detention may be the initial step in a (C) ……………..copy of the saga of the (Au)
………………. Dust Kid.
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Laboratory Techniques for Environmental Technicians
Practical work
Complete the following table by observing samples of elements that you may have access to
at home
What criteria did you use to decide that an element was a metal or a non-metal?
Type you answer here
Leave blank for assessor feedback
Compounds are substances composed of two or more atoms chemically combined in a fixed
ratio. Elements are always present in the same ratio in a given compound. The properties of
a compound are usually quite different from those of the elements from which it is
composed. Chemical methods are required to separate compounds into their constituent
elements.
List any compounds that you are familiar with eg water
Type you answer here
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A mixture is a physical combination of two or more substances. A mixture has a variable
composition and may be identified as heterogeneous or homogeneous. Mixtures can be
separated into pure substances using methods based on differences in their physical
properties.
Examples
a magnet can be used to separate iron from other solids……magnetic properties
solids and liquids can be separated by decanting, filtering and centrifuging……. differences in
density or particle size
two immiscible liquids can be separated using a separating funnel
List some mixtures that you are familiar with.
Type you answer here
Leave blank for assessor feedback
Homogeneous mixtures (solutions) have uniform properties throughout. Solutions may be
gases, liquids or solids. Like all other mixtures solutions have variable composition.
Heterogeneous mixtures are those that do not have uniform properties throughout.
Mixtures can be easily separated:
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Laboratory Techniques for Environmental Technicians
Laboratory techniques for separating mixtures
METHOD
MIXTURE TYPE
DIAGRAM
(property differences)
Decanting
Used
to
undissolved
separate
sediment from a liquid
(density)
or
immiscible liquids.
Separating funnel
used
to
separate
immiscible liquids
(density)
Filtration
Gravity and
techniques
used to separate an
Vacuum undissolved solid from a
liquid.
(particle size)
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Laboratory Techniques for Environmental Technicians
METHOD
MIXTURE TYPE
DIAGRAM
(property differences)

used
to
separate
undissolved
substances
(some liquid) of differing
densities
Centrifuge
(density)
Evaporation
Crystallisation
/  used to separate
dissolved solid from
solvent (liquid)
Distillation
a
a
 used to separate mixtures of liquids
(boiling point)
Chromatography
 used to separate miscible
liquids
(partition – differences in
attraction of sample
matrix for mobile phase
Vs stationary phase)
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Laboratory Techniques for Environmental Technicians
METHOD
MIXTURE TYPE
DIAGRAM
(property differences)
Which laboratory method would you use to separate the following mixtures?
◗ charcoal and salt
◗ water and oil
◗ alcohol and water
◗ sugar from a sugar solution
◗ water from muddy river water
Type you answer here
Leave blank for assessor feedback
Classify each of the following as element, compound or mixture
◗ soil
◗ silver
◗ milk
◗ table sugar
◗ sulfur
◗ river water
◗ grape juice
◗ nitrogen
◗ fog
Type you answer here
Leave blank for assessor feedback
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Laboratory Techniques for Environmental Technicians
A change in the properties of a substance without a change in composition is a physical
change.
Physical changes include melting and boiling eg water to ice and steam
Changes of State
We can show changes of state in the following way:
If there is a change in the composition of a substance, however, a chemical change is
indicated. In a chemical change reactants are converted to products eg the breakdown of
water into its elements hydrogen and oxygen.
In any physical or chemical change, both mass and energy are conserved.
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Laboratory Techniques for Environmental Technicians
Practical: Physical versus chemical change
Purpose: To distinguish a physical from a chemical change based on observation.
Task: (Teacher demonstration)
◗ You will be required to record observations for each set up and identify each as either a
physical or chemical change.
◗ Summarise the process(es) taking place at each station, expressing each in terms of
starting and finishing substances. Expressing physical states as subscripts next to each
substance will assist in this regard.
◗ Describe the demonstrations of the reactions. State whether a chemical or physical
change takes place.
o Zinc and acid
o Copper and acid
o Sodium and water
o Sulfur, air and a flame
o Magnesium, air and a flame
o Iodine and heat
o Boiling water
o Electrolysis of water
o Hydrogen, air and flame
Question
What observations could you use to assist in determining if a chemical or physical change is
occurring?
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Laboratory Techniques for Environmental Technicians
Atomic Structure
Plato, Aristotle
500 BC
All matter composed of four elements, earth,
wind, air and fire
Democritus
400BC
Matter not infinitely divisible, but was made of
extremely small hard particles called atoms
which were indivisible
Cavendish
1700’s
Made water from hydrogen and oxygen 
disproving previous theory
Lavoisier
1700’s
Law of conservation of mass: matter cannot be
created or destroyed
Proust
1799
Law of constant composition: a compound
always contains the same elements in the
same ratios by mass
Dalton
1808
1. Matter is composed of tiny indivisible
particles called atoms
2. All atoms of the one element are the same
but different to atoms of other elements
3. Chemical reactions consist of combining ,
separating or rearranging atoms in simple
whole number ratios
Thomson
1904
Discovery of electrons as particles within
atoms. Proposed “plum pudding” model,
where electrons were embedded in a sphere
of positive charge
Rutherford
1911>
Famous gold leaf experiment:
1. Matter is mostly empty space
2. The atoms consists of a small dense +ve
nucleus containing most of the mass of the
atom
3. The nucleus is surrounded by electrons
making it electrically neutral
Chadwick
1932
Discovered the existence of neutrons
Bohr
1913>
Proposed models of atom which contained
allowable orbits or shells for electrons
surrounding the nucleus.
Many other contributions have been made by famous scientists such as Plank, Schrodinger,
Oppenheimer etc. The model of the atom is being constantly refined.
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Laboratory Techniques for Environmental Technicians
Bohr model of the atom
Proposed by Neils Bohr;
◗ central massive body, nucleus, (like the sun in our solar system)
◗ The nucleus has a positive charge
◗ very small electrons circling, (like the planets)
Terms related to atomic structure
Term
Definition
Atomic number (Z)
the number of protons in an atom. Identifies the element
Atomic weight
average mass of naturally occurring isotopes of an element
Electrons
negatively charged particles orbiting the nucleus
Ions
a charged atom
Isotopes
atoms of the same element with differing numbers of neutrons
Mass number (A)
the number of protons plus the number of neutrons in an
atom
Neutrons
particles found in the nucleus but with no charge
Protons
positively charged particles found in the nucleus
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Laboratory Techniques for Environmental Technicians
An element may be represented in the Periodic Table using a shorthand notation: ZXA
Where:
X
=
element symbol
Z
=
atomic number, number of protons
A
=
mass number
Neutron number
= mass number - atomic number
=
A-Z
Particle
Relative charge
Relative mass
Location
proton
+1
1 au
Inside the nucleus
neutron
0
1 au
Inside the nucleus
Electron
-1
0
Orbiting the nucleus
Use your periodic table to complete the following table, remembering the following
information
Atomic number = identity of atom = number of protons = no of electrons(for an uncharged
atom)
Mass number = no of protons + number of neutrons
Neutrons = Mass number – number of protons
Element
Atomic
(Z)
hydrogen
1
1
Carbon
6
Text
Text
Text
Text
Text
26
Text
Text
Text
Text
Text
Text
No. Mass No. (A) Protons (p)
Electrons (é)
Neutrons (n)
1
1
0
13
6
6
7
Text
Text
11
12
22
11
Text
Text
Text
Text
Text
Text
210
197
Text
Text
Text
36
Text
79
29
48
128
Text
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Laboratory Techniques for Environmental Technicians
How are the electrons arranged around the nucleus?
For this study we will use only the simplified version and only look at the first 20 elements.
Each shell (or orbit) about the nucleus has a specific number of electrons that can be held. In
electron, in an unexcited atom, ground state) can only be placed in a shell in a certain order.
A shell must be filled before a new shell can commence. The pattern for filling the shells
follows the following pattern:
◗ Shell 1 can contain a maximum of 2 electrons
◗ Shell 2 can contain a maximum of 8 electrons
◗ Shell 3 can contain a maximum of 8 electrons
◗ Shell 4 can contain a maximum of 18 electrons
Electron configuration of the first 20 elements on the Periodic Table
Atomic
No
Element
Protons
Electrons
1
Hydrogen
1
1
(found
shell)
2
Helium
2
Lithium
3
in
1st
in
2nd
e
2
(found
shell
3
Shell diagram
e
3
e
Text
(2 electrons in
1st shell +
1 electron in 2nd
shell)
4
Text
Text
Text
Text
5
Text
Text
Text
Text
6
Text
Text
Text
Text
7
Text
Text
Text
Text
8
Text
Text
Text
Text
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Laboratory Techniques for Environmental Technicians
9
Text
Text
Text
Text
Text
10
10
Text
10
11
(2 electrons in
1st shell + 8
electrons in 2nd
shell, which is
now full)
Sodium
11
11
Text
(2 electrons in
1st shell + 8
electrons in 2nd
shell
+
1
electron in 3rd
shell)
12
Text
Text
Text
Text
13
Text
Text
Text
Text
14
Text
Text
Text
Text
15
Text
Text
Text
Text
Text
Text
Text
Text
17
Text
Text
Text
Text
18
Text
Text
Text
Text
19
Text
19
19
Text
16
(2 electrons in
1st shell + 8
electrons in 2nd
shell
+
8
electrons in 3rd
shell
+
1
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Laboratory Techniques for Environmental Technicians
electron
4thshell)
20
Text
Text
in
Text
Text
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Laboratory Techniques for Environmental Technicians
How can atoms become charged?
In simple terms the only part of an atom that is able to move is the electron. Electrons can
be added or removed from an atom. If this occurs then the atom is no longer neutral and a
charged particle, called an ion is formed.
If electrons are lost the atom becomes positively charged and is termed a cation
If electrons are gained the atom becomes negatively charged and is termed an anion.
For example.
Sodium atom:
◗ has 11 protons and a nuclear charge
= +11
◗ has 11 electrons and an electron charge
= -11
◗ overall net charge
=
0
Sodium loses an electron easily forming a sodium cation
◗ 11 protons and a nuclear charge
= +11
◗ Now has 10 electrons and an electron charge
= -10
◗ Overall net charge
= +1
Oxygen atom:
◗ Has 8 protons and a nuclear charge
= +8
◗ Has 8 electrons and an electron charge
= -8
◗ Overall net charge
= 0
Oxygen can gain two electrons easily forming a oxygen anion
◗ Has 8 protons and a nuclear charge
= +8
◗ If it gains 2 electrons the electron charge
= -10
◗ Overall net charge
= -2
Explain how the following can occur
Explanation
Text
Chlorine has -1 charge
Text
Magnesium has +2 charge
Text
Phosphorus has -3 charge
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Laboratory Techniques for Environmental Technicians
Text
Aluminium has + 3 charge
Text
Sulfur has -2 charge
Text
Caesium has +1 charge
Is there a pattern you can see in the relative charges of the metals and non-metals?
Type you answer here
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Laboratory Techniques for Environmental Technicians
Terms and Definitions (From Chemistry for Technicians)
Terms
A
proton
B
electron
C
neutron
D
atomic number
E
isotope
F
mass number
G
atomic weight
H
ion
I
anion
J
cation
K
monatomic ion
L
nucleus
M
valence shell
Definitions
1
The number of protons in an atom
Term
2
A negatively charged ion
Type your answer here
3
A positively charged subatomic particle
Type your answer here
4
Atoms with the same number of protons, Type your answer here
but different numbers of neutrons
5
A charge atom or molecule
Type your answer here
6
An ion containing a single atom
Type your answer here
7
The centre of the atom, which contains Type your answer here
most of the mass
8
A negatively charged subatomic particle
Type your answer here
9
A positively charged ion
Type your answer here
10
The combined total of protons and neutron
Type your answer here
11
The outermost shell containing electrons
Type your answer here
12
A subatomic particle with no charge
Type your answer here
13
The average mass of naturally occurring Type your answer here
atoms of an element
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Laboratory Techniques for Environmental Technicians
The Periodic Table & Atomic Structure
The periodic table is an arrangement of the known elements. The modern periodic table has
the elements arranged in order of increasing atomic number…….the early versions had the
elements arranged according to atomic mass. Can you find examples where this would be
different to the modern version?
The periodic table briefly
Independently arrived at by Mendeleev and Meyer in the mid-19th century. Mendeleev
predicted the properties of undiscovered elements and left spaces in his periodic table.
Consists of a series of horizontal rows, periods and vertical columns, groups:
Group 1:
Alkali metals
Group 2:
Alkaline earth metals
Group 7:
Halogens
Group 8:
Inert gases
provides a summary of atomic weights and numbers, symbols and names for the elements
understanding of properties of related elements
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Laboratory Techniques for Environmental Technicians
Chemical Formulae
A chemical symbol represents one atom of an element. A chemical formula represents one
molecule or one formula unit of an element or compound The formula shows:
◗ Each type of atom present
◗ The number of each atom present in the molecule or formula unit
◗ Coefficients, superscripts and subscripts where necessary
A molecule is a group of atoms which are bonded together…the atoms may be the same or
different.
Examples:
Elements
Compounds
H2 = hydrogen
H2O
P4 = phosphorus
C6H12O6 = glucose
O2 = oxygen
NH3
= water
= ammonia
A formula unit represents the simplest ratio of atoms in a substance, which may be a giant
lattice
Examples:
Substance
Type
Formula
Diamond
Element
C
Sodium chloride
Compound (ionic)
NaCl
Silica
Compound (covalent)
SiO2
Ions (charged particles) use superscripts to show the electrical charge.
Example
Mg2+ = a magnesium ion having a charge of +2
CO32- = one carbon atom and three oxygen atoms forming an ion, carbonate ion, with a
charge of –2
 For each of the following describe the information each formula is providing about the
elements and how many of them there are.
CH4
= 1 carbon atom and 4 hydrogen atoms
Sn3(PO4)2
=
CuSO4.5H2O
=
Cl-
=
OH-
=
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Laboratory Techniques for Environmental Technicians
The elements and their symbols (arranged alphabetically by symbol)
Ac
Ag
Al
Am
Ar
As
At
Au
B
Ba
Be
Bi
Bk
Br
C
Ca
Cd
Ce
Cf
Cl
Cm
Co
Cr
Cs
Cu
Dy
Er
Es
Eu
F
Fe
Fm
Fm
Ga
Gd
Ge
H
He
Hf
Hg
Ho
actinium
silver
aluminium
americium
argon
arsenic
astatine
gold
boron
barium
beryllium
bismuth
berkelium
bromine
carbon
calcium
cadmium
cerium
californium
chlorine
curium
cobalt
chromium
caesium
copper
dysprosiu
m
erbium
einsteiniu
m
europium
fluorine
iron
fermium
francium
gallium
gadolinium
germanium
hydrogen
helium
hafnium
mercury
holmium
I
In
Ir
K
Kr
La
Li
Lr
Lu
Md
Mg
Mn
Mo
N
Na
Nb
Nd
Ne
Ni
No
Np
O
Os
P
Pa
Pb
iodine
indium
iridium
potassium
krypton
lanthanum
lithium
lawrencium
lutetium
mendelevium
magnesium
manganese
molybdenum
nitrogen
sodium
niobium
neodymium
neon
nickel
nobelium
neptunium
oxygen
osmium
phosphorus
proactinium
lead
Pd
P
m
Po
Pr
Pt
Pu
Ra
Rb
Re
Rh
Rn
Ru
S
Sb
Sc
palladium
promethium
Se
Si
Sm
Sn
Sr
Ta
Tb
Tc
Te
Th
Ti
Tl
U
Unq
Unp
Unh
V
W
Xe
Y
Yb
Zn
Zr
selenium
silicon
samarium
tin
strontium
tantalum
terbium
technetium
tellurium
thorium
titanium
thallium
uranium
104 in dispute
105 in dispute
106 in dispute
vanadium
tungsten
xenon
yttrium
Ytterbium
Zinc
Zirconium
polonium
praseodymium
platinum
plutonium
radium
rubidium
rhenium
rhodium
radon
ruthenium
sulfur
antimony
scandium
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Laboratory Techniques for Environmental Technicians
Systematic naming of inorganic compounds
Inorganic compounds are generally of two types:
◗ Ionic (generally composed of a metal ion and a non-metallic ion eg sodium chloride,
NaCl)
◗ Molecular (generally composed only of non-metals eg nitrogen dioxide, NO2)
To manipulate formulae, it is necessary to:
◗ learn the symbols for the common elements (or be able to obtain them from the
periodic table)
◗ learn the common ions and their charges
◗ identify types of formulae. learn to name compounds correctly.
Common ions and their charges
+1
+2
+3
+4
-1
-2
-3
ammonium
NH4+
barium
Ba2+
aluminium
Al3+
phosphate
PO43-
calcium
Ca2+
Copper (II)
Cu2+
iron (II)
Fe2+
lead (II)
Pb2+
magnesium
Mg2+
iron (III)
Fe3+
acetate
(ethanonate)
CH3COO bromide
Br chlorate
ClO3 chloride
Cl fluoride
Fhydrogen
carbonate
HCO3hydrogen sulfate
HSO4 hydroxide
OH iodide
Initrate
NO3 nitrite
NO2 Permanganate
MnO4-
carbonate
CO32-
potassium
K+
silver
Ag+
sodium
Na+
Hydrogen
H+
Lead
(IV)
Pb4+
tin (IV)
Sn4+
chromate
CrO42dichromate
Cr2O72oxide
O2peroxide
O22sulfate
SO42-
Phosphide
P3Nitride
N3-
mercury(II)
Hg2+
nickel
Ni2+
tin (II)
Sn2+
sulfite
SO32sulfide
S2-
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Laboratory Techniques for Environmental Technicians
Types of compounds
The simplest compounds contain only two elements and are called binary compounds.
a) Metal-non-metal binary compounds (ionic compounds)
The metal is written as the first name and the non-metal ending “ide” forms the last name
eg Na2S sodium sulfide
MgO magnesium oxide
If more than one compound of the metal exists, their oxidation numbers are used
eg FeCl2 iron (II) chloride
and
FeCl 3 iron (III) chloride
Name the following metal-non-metal compounds (check for a metal then it is as simple as
writing the metal name followed by the non-metal ending.
ZnO
Type your answer here
K2O
Type your answer here
Al2O3
Type your answer here
Ca3(PO4)2
Type your answer here
Complete the table below by writing in the name of the metal and the non-metal ending
then combining them together
Ions
Li+
NH4+
Mg2+
Lithium
Br -
Al3+
Na+
Aluminium
Text
Text
Text
Text
Bromide
Lithium
bromide
SO42-
Text
Text
Text
Text
Text
OH -
Text
Text
Text
Aluminium
hydroxide
Text
F-
Text
Text
Text
Text
Text
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Laboratory Techniques for Environmental Technicians
O2-
Text
Text
Text
Text
Text
NO3 -
Text
Text
Text
Text
Text
PO43-
Text
Text
Text
Text
Text
Cl -
Text
Text
Text
Text
Text
S2-
Text
Text
Text
Text
Text
I-
Text
Text
Text
Text
Text
CO32-
Text
Text
Text
Text
Text
HCO3 -
Text
Text
Text
Text
Text
CH3COO-
Text
Ammonium
acetate
Text
Text
Text
(ammonium
ethanoate)
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Laboratory Techniques for Environmental Technicians
Chemical formulae of ionic compounds
The valency of an atom or polyatomic ion can be viewed as a bonding position which must
be filled to form a stable compound.
Recap: in order to find the formula of a compound you must recognise the parts or
constituents from the name. Since the name is in two parts this recognition is easy as long
as you realise that the ending of the name of the second atom is often changed to -ide.
Oxygen becomes oxide; sulfur becomes sulfide; chlorine becomes chloride etc.
Na
a
Na
Cl
Cl
will join the ratio 1:1
 Sodium’s valency is 1+
Formula is Na1Cl1 or
 Chloride’s valency is 1-
NaCl (1 is automatically assumed)
2.
Aluminium nitrate is a compound made from aluminium and nitrate polyatomic ion
Al
NO3
NO3
Al
will join in the ratio 1:3
NO3
Aluminium has a valency of 3+
NO3
Nitrate has a valency of 1Formula will be Al(NO3)3
3.
Lead IV oxide is a compound made from lead (valency 4) and oxygen.
Pb
O
will join in the ratio 1:2
Pb
O
Lead has a valency of 4+
Oxygen has a valency of 2The formula will be PbO2
O
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Laboratory Techniques for Environmental Technicians
Cross Multiply Method
Step 1. Write the symbols which represent each part of the name.
Step 2. Write the valencies as a superscript to each part.
Step 3. Cancel by dividing the valencies by any common factor
Step 4. Cross over the numbers to form subscripts, using brackets where necessary.
Examples:
1+
1.
Na
Sodium Chloride
1-
Cl
3+
2.
Aluminium Nitrate
Al
1-
NO3
4+
3.
Lead (IV) Oxide
Pb
2 1
Formula is Na1Cl1 or NaCl
Formula is Al1(NO3)3 or Al(NO3)3
2-
O
Formula is Pb1O2 or PbO2
4 2
Write chemical formulae for the following compounds.
Compound
Formula
Compound
Formula
calcium chloride
Text
potassium sulfite
Text
magnesium oxide
Text
nickel carbonate
Text
lead (IV) sulfate
Text
silver phosphate
Text
Text
potassium nitrate
Text
Text
sodium carbonate
Text
Text
ammonium chloride
Text
barium phosphide
zinc sulfate
iron (III) hydroxide
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Laboratory Techniques for Environmental Technicians
b) Two non-metal binary compounds (molecular compounds)
The first name is the more electropositive element and the other non-metal is given the -ide
ending for the surname.
The prefixes:
◗ mon
=1
◗ di = 2
◗ tri = 3
◗ tetra
=4
◗ penta
=5
◗ hexa
=6
are used to show the proportions of each element. eg
N2O3
dinitrogen trioxide
SO2
sulfur dioxide
The prefix mono is used only when there is more than one compound of the two elements
eg
CO
carbon monoxide
CO2
carbon dioxide
Write the formula for the following non-metal binary compounds
Compound
Formula
Compound
Formula
nitrogen triiodide
Text
silicon dioxide
Text
silicon tetrafluoride
Text
water
Text
Diphosphorus
pentoxide
Text
ammonia
Text
c) Common names are given to some hydrides
H2O
water
NH3
ammonia
CH4
methane
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Laboratory Techniques for Environmental Technicians
d) Composite ion (radical) compounds
The positive ion is written as the first name and the negative ions the surname. eg
NH4Br
ammonium bromide
NaOH
sodium hydroxide
(NH4)2SO4
ammonium sulfate
e) Acids
i) Halogen acids eg HCl(g)
HCl(aq)
hydrogen chloride
hydrochloric acid
ii) Oxygen containing acids
If only one acid exists, the ending -ic is used eg
hydrofluoric acid
HF
carbonic acid
H2CO3
If two acids exist the lowest oxidation state ending is -ous and the higher state ending is -ic.
eg
sulfurous acid
H2SO3
sulfuric acid
H2SO4
The formula for the some common acids is provided below. (In some cases the common
name has also been given).
Name
Common name
Formula
Hydrochloric acid
Muriatic acid
HCl
Sulfuric acid
Oil of vitriol
H2SO4
Phosphoric acid
Text
H3PO4
Nitric acid
Text
HNO3
Sulfurous acid
Text
Text
Ethanoic acid
Acetic acid
Text
Boric acid
Text
Text
Carbonic acid
Text
Text
Phosphoric acid
Text
Text
Name the following compounds (Remember to identify the type of compound first eg is
there a metal …… if yes then just name the metal and the ending!)
Formula
Name
Formula
Name
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Ni(NO3)2
Text
SO3
Text
PCl3
Text
FeSO4
Text
P2O5
Text
Fe2(SO4)3
Text
Fe2O3
Text
PbO
Text
CrCl3
Text
PbO2
Text
IF7
Text
Ca3N2
Text
KHSO4
Text
NO2
Text
Li3PO4
Text
N2O4
Text
NaNO2
Text
Ca(OH)2
Text
LiH
Text
Na2CO3
Text
S2Cl2
Text
Text
Complete the following table
Name
Formula
Name
Formula
copper (II) oxide
Text
Ammonia
Text
oxygen difluoride
Text
potassium permanganate
Text
copper (II) sulfate
Text
hydrogen chloride
Text
barium hydroxide
Text
dihydrogen sulfide
Text
copper (I) bromide
Text
Silicon dioxide
Text
aluminium hydroxide
Text
phosphorus
pentachloride
Text
dinitrogen tetroxide
Text
chromium (III) fluoride
Text
Silver bromide
Text
Text
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Laboratory Techniques for Environmental Technicians
Organic chemicals
Organic chemistry involves the study of carbon compounds which contain hydrogen and
sometimes oxygen, chlorine, nitrogen etc. Many of out useful organic chemicals are
obtained from coal and oil or synthesised from chemicals obtained from them or plants and
animals.
Coal    benzene, toluene, naphthalene, creosote
Petroleum   methane, propane, octane, petrols, oils, kerosene, paraffin, bitumen
Plants    sugar, starch, edible oils, waxes, gelatin, dyes, drugs, natural fibres
These raw materials are either used directly or converted to others eg plastics,
pharmaceuticals, dyes, pigments, cosmetics, insecticides, explosives, refrigerants, paints etc
Carbon forms so many compounds because of its ability to form strong bonds to itself and
other elements. IT ALWAYS HAS 4 BONDS (places where other elements can join)
The other elements carbon bonds with include;
◗ Hydrogen
1 bond
◗ Oxygen
2 bonds
◗ Nitrogen
3 bonds
◗ Chlorine
1 bond
◗ Bromine
1 bond
◗ Iodine
1 bond
Carbon can also form multiple bonds with itself and oxygen and nitrogen
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Laboratory Techniques for Environmental Technicians
Carbon compounds are named according to the family they belong to according to their
structure.
Family
Structure contains
Alkanes
Single bonds between C Methane
atoms
Butane
Alkenes
Double bond
carbons
between Ethene
Triple
bond
carbons
between Ethyne
Alkynes
Alkanols (alcohols)
Example
Butene
Contain –OH
Butyne
Methanol
Butanol
Alkanals (aldehydes)
Contain C=O on the end
Methanal
Butanal
Alkanones (ketones)
C=O within the structure
Propanone
Butanone
Alkanoic acids
Methanoic acid
Butanoic acid
Esters
Methyl butanoate
Butyl methanoate
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Laboratory Techniques for Environmental Technicians
Writing Chemical Equations - Chemical Shorthand
A chemical equation is a shorthand method of communicating information about reactants
and products, quantities involved and sometimes how much energy is needed or released. It
accounts for all the atoms involved in the re-arrangement of bonds.
During a chemical reaction, bonds are broken in the reactants and new bonds are formed
making the products. The atoms are re-arranged during a reaction, but the numbers of the
different types of atoms remain constant.
The law of Conservation of Matter (or mass) states that during a chemical reaction the total
mass of reactants is the same as the total mass of products formed. During a chemical
reaction, matter is neither created nor destroyed, it is only changed from one form into
another.
Law of Conservation of Mass (Matter)
Note: The total number of atoms of each element is unchanged!
Example 1:
When a spark ignites a mixture of hydrogen gas and oxygen gas, an explosion
occurs - a large amount of energy is released. The product formed is water. Chemical
reactions which release energy are called exothermic.
Word Equation:
Hydrogen(g) + Oxygen(g) ⇌ Water(l) + energy
Symbol Equation:
H2(g) + O2(g) ⇌ H2O(l) + energy (unbalanced)
The correct chemical formula is written under each reactant and product. The extra
subscripts indicate the state of that substance.
The equation above is said to be unbalanced because the numbers of various atoms
represented on the reactants’ side (left of the arrow) is not the same as those on the
products’ side (right of the arrow). This would mean that mass is not conserved!
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Laboratory Techniques for Environmental Technicians
Reactants ( H2 + O2 )
Products ( H2O )
H=2
H=2
O=2
O=1
The numbers of oxygen atoms are not balanced!
We must adjust the numbers of molecules of each type present to reach a balance! This
involves placing numbers where required in front of each formula. The number will multiply
each atom in the formula. This was referred to previously as a coefficient number. A “2” in
front of the H2O product will provide the required 2 oxygen atoms on the right.
H2(g) + O2(g) ⇌ 2H2O(l) + energy
Reactants ( H2 + O2 )
Products ( 2 H2O )
H=2
H=4
O=2
O=2
Whilst oxygen has now been balanced, Hydrogen is now unbalanced!
In order to balance the hydrogen atoms it is necessary to place a “2” in front of the H2 on
the left. This provides 4 hydrogen atoms on the left also.
Balanced Equation: 2H2(g) + O2(g) ⇌ 2H2O(l) + energy
Reactants ( 2 H2 + O2 )
Products ( 2 H2O )
H=4
H=4
O=2
O=2
Example 2
When an emergency flare lights up, magnesium (or aluminium) metal combines with
oxygen gas releasing a bright light. This is also an exothermic reaction. The
substance formed is a metal oxide.
Word equation:
Magnesium(s) + Oxygen(g) ⇌ Magnesium oxide(s) + energy
Symbols:
Mg(s)
+
O2(g)
⇌
MgO(s)
+
energy
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Reactants ( Mg + O2 )
Products ( MgO )
Mg = 1
Mg = 1
O=2
O=1
The numbers of oxygen atoms are not balanced!
Balance the oxygen atoms by placing a coefficient “2” in front of MgO. This will also
produce 2 magnesium atoms.
Mg(s) + O2(g) ⇌ 2MgO(s) + energy
Reactants ( Mg + O2 )
Products ( 2 MgO )
Mg = 1
Mg = 2
O=2
O=2
Then balance the magnesium atoms by placing a coefficient “2” in front of Mg.
Balanced equation: 2 Mg(s) + O2(g) ⇌ 2MgO(s) + energy
Reactants ( 2 Mg + O2 )
Products ( 2 MgO )
Mg = 2
Mg = 2
O=2
O=2
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Laboratory Techniques for Environmental Technicians
Written Exercises
1.
Balance the following chemical equations.
(a)
CH4(g) +
O2(g)
⇌
CO2(g)
+
H2O(l)
(b)
Na(s)
+
H2O(l )
⇌
NaOH(aq)
+
H2(g)
(c)
Zn(s)
+
HCl(aq )
⇌
ZnCl2(aq)
+
H2(g)
(d)
P
+
O2
⇌
P2O5
(e)
NH3
+
H2SO4
⇌
(NH4)2SO4
(f)
CuO
+
HCl
⇌
CuCl2
+
H2O
(g)
H2O2
⇌
H2O
+
O2
(h)
H2CO3
⇌
H2O
+
CO2
(i)
Fe
+
O2
⇌
Fe2O3
(j)
C8H18 +
O2
⇌
CO2
+
H2O
The Mole Concept
We need to “scale up” our quantities from the molecular level to an amount which we can
see and measure in the laboratory. The quantity of particles chosen is called the mole which
is approximately 6.02 x 1023 particles of any pure substance (ie; element or compound).
Scientific notation is expressed in this number as this is a gigantic quantity!
The Periodic Table measures the mass of each element in atomic units (a.u.)
One mole of any substance contains 6.02 x 1023 particles and has a mass measured in grams
equal to its formula weight (a.u.)
One mole of water, H2O contains 6.02 x 1023 molecules. The mass of 1 mole of water is
equal to its formula weight (FW), expressed in grams;
FW (Hydrogen) = 1.008 a.m.u.
FW (Oxygen) = 16.00 a.m.u.
∴ (2 x 1.008) + 16.00 = 18.016 g.
Likewise, one mole of hydrogen (gas) contains 6.02 x 1023 molecules of hydrogen and has a
mass of 2 x 1.008 = 2.016 g.
And one mole of oxygen (gas) contains 6.02 x 1023 molecules of oxygen and has a mass of 2
x 16.00 = 32.00 g.
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Laboratory Techniques for Environmental Technicians
We could, if we wanted, determine the ratio, by mass of the substances involved in the
decomposition of water.
⇌
2H2O(l)
2 molecules
2H2(g)
+
2 molecules
of water
O2(g)
1 molecule
of hydrogen
of oxygen
2 moles
2 moles
1 mole
of water
of hydrogen
of oxygen
2 x 18.016
2 x 2.016
1 x 32.00
= 36.032 g
= 4.032 g
= 32.00 g
Therefore;
If we started with 1 mole (ie 36.032 g) of water, we would obtain 4.032 g of hydrogen and
32.00 g of oxygen. What mass of hydrogen and oxygen would we obtain if we started with
1.00 g of water?
H2O (l)
H2 (g)
O2 (g)
36.032
4.032
32.00
36.032
36.032
36.032
Therefore
1.00 g
⇌
0.11 g
+
0.89 g
If we started with 1.00 g of water we would obtain 0.11 g of hydrogen and 0.89 g of oxygen.
Notice that the balancing numbers do not give us masses directly, that is we cannot say that
2 g of water will give 2 g of hydrogen and 1 g of oxygen!
When we use the mole concept we are taking into account that the various atoms have
different masses, and the balancing numbers give us the right ratios of atoms involved.
It makes life much easier when interpreting number coefficients of a balanced
chemical equation as mole quantities, because that’s exactly what they are !
Take special note that the total mass of reactants equals the total mass of products.
No atoms are lost or gained, as understood from our work on writing and balancing
chemical equations. Atoms are just re-arranged in a chemical reaction.
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Laboratory Techniques for Environmental Technicians
We could continue our calculations to account for any mass of water, or we could calculate
the quantity of water required to produce a particular amount of hydrogen or oxygen. The
mole ratios of reactants to products are determined from the balanced chemical equation
and a knowledge of the mole concept.
We can determine the number of moles present in a sample of a substance by using one or
more of the following formulae:
No of moles
=
No of particles
6.02 x 1023
This mole equation will not be immediately useful as it deals with a known number of
chemical particles, which in reality is mostly unrealistic. You may, however be asked to
calculate a mole quantity given such data.
No of moles
Mole Equation 1.
=
Mass (g)
Formula Weight
This equation is useful in the real world, as mass and Formula Weight are measurable
quantities in the real world.
Shorthand:
n = m
FW
Written Exercises:
Mole calculations
1. Calculate the Formula Weight, (FW) for each of the following substances for which the
formula is provided.
(a) Sodium Chloride, NaCl
Type your answer here
(b) Ammonium Sulfate, (NH4)2SO4
Type your answer here
(c) Glucose, C6H12O6
Type your answer here
(d) Potassium Hydrogen Phthalate, KC8H5O4
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Laboratory Techniques for Environmental Technicians
Type your answer here
2. Calculate the number of moles of each substance
(a) 58 g of NaCl
Type your answer here
(b) 13.2 g of (NH4)2SO4
Type your answer here
(c) 90 g of C6H12O6
Type your answer here
(d) 5.0 g of KC8H5O4
Type your answer here
3. What mass of each substance is present?
(a) 2.5 moles of NaCl
Type your answer here
(b) 0.5 moles of (NH4)2SO4
Type your answer here
(c) 4 moles of C6H12O6
Type your answer here
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Laboratory Techniques for Environmental Technicians
Solutions
When a substance dissolves in a liquid we say a solution has been formed . The substance
can be a solid, liquid or gas and is called the solute. The liquid it dissolves in is called the
solvent. The strength or concentration of a solution can be described by the number of
moles of solute in 1 litre of solution (other units may be % or g per litre). An aqueous
solution is the term applied for a solution made of a solute with water as the solvent.
A 1 molar solution (1M) has 1 mole of solute dissolved in 1 litre of solution
Complete the following:
A 0.5 molar sucrose solution (0.5M) has 0.5 moles of sucrose in 1 litre of solution.
A 2.5 M salt solution hasClick here to enter text.moles of salt in 1 litre of solution
A 0.001 M NaOH solution hasClick here to enter text. moles of Click here to enter text.in 1
litre of solution
A 0.2 M KI solution has Click here to enter text.in 1 litre of solution.
No. of moles of solute = (Concentration of solution) X (Volume of solution)
(M)
X
(L)
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Laboratory Techniques for Environmental Technicians
Calculating the molarity of a solution
The number of moles (n) of solute per litre (L) of solvent is referred to as the MOLARITY of a
solution. Molarity can be written as mol/L, but it is more commonly written as M. Hence
when you are given a bottle marked ‘0.1M NaOH’ this means that the solution consists of
0.1 mole of sodium hydroxide per litre of solvent. This solution is referred to as being a 0.1
Molar solution. [ ] represents concentration in moles/L
Molarity = moles
(moles/L)
litres
Examples
1.What is the molarity of a solution containing 0.25 moles of NaOH in 1.2 L of water.
[ NaOH] = 0.25 moles
1.2 L
= 0.208 mol/L
2. How many grams of NaOH are required to make 2 litres of a 0.15 M solution?
a. Firstly calculate how many moles of NaOH are required to make up the solution.
Rearrange the equation to make ‘moles’ the subject of the formula
moles (n) = Molarity (M) x Litres (L)
moles
= 0.15 x 2
moles
= 0.3
b. How many grams of NaOH are there in 0.3 mole of the substance?
moles
= mass (g)
FWt
Rearrange equation to make mass the subject.
mass
= moles x FWt
Calculate the formula weight of NaOH = 39.998
Mass
= 0.3 x 39.998
Mass
= 11.999g NaOH
Therefore, 11.999g NaOH are required to prepare 2 L of 0.15 M NaOH
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Laboratory Techniques for Environmental Technicians
Exercises
1. Determine the molarity of a solution in which 4.67 moles of Li2SO3 are dissolved to make
2 L of solution
Type your answer here
2. Determine the molarity of a solution in which 4.783 g of Na2CO3 is dissolved to make 10 L
of solution.
Type your answer here
3. Determine the number of moles of solute needed to prepare 16.00 mL of a 0.415 M
Pb(NO3)2 solution.
Type your answer here
4. Determine the mass of solute required to prepare 0.500L of a 1.00 M KCl solution.
Type your answer here
5. Determine the final volume of the solution needed when 8.07g of (NH4)2CO3 is dissolved
to make a 0.250 M solution.
Type your answer here
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Laboratory Techniques for Environmental Technicians
Dilution of solutions
Many laboratories will keep a range of commonly used solutions, often at a fairly high
Molarity. One of your tasks may be to prepare a solution of lower Molarity from one of
these ‘stock’ solutions. You will be diluting the stock solution. To dilute a solution means to
add more solvent without the addition of more solute.
There is an equation which relates the Molarity and Volume of the stock solution to the
Molarity and Volume of the solution you wish to prepare.
C1V1 = C2V2
where C = concentration, which is expressed as molarity or moles/L
V = volume which is expressed in litres
C1 refers to the concentration of the stock or ‘initial’ solution, and V 1 refers to the
volume of the initial solution.
C2 refers to the concentration of the final solution, and V 2 refers to the volume of the
final solution.
Example
Using a 2M stock solution of NaOH prepare 500 mL of a 0.125 M solution
C1 = 2M
C2 = 0.125 M
V1 = ?
V2 = 500 mL
Rearrange the formula to make V1 the subject
V1 = C2V2
ie V1 = 0.125 M x 500mL
C1
2M
V1 = 31.3 mL
Hence if 31.3 mL of 2 M NaOH is diluted to 500 mL a solution of 0.125M will be made.
The units on both sides of the equation must be the same.
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Laboratory Techniques for Environmental Technicians
Exercises
1. A stock solution of 0.750 M NaCl is available. What volume, in mL, is required to make
100.0 mL of 0.10 M?
Type your answer here
2. Concentrated H2SO4 is 18.0 M. What volume is required to prepare 2.00L of a 1.00M
solution?
Type your answer here
3. Concentrated HCl is approximately 11.8 M. What volume is required to prepare 2.0L of a
1.0M solution?
Type your answer here
4. A 0.500 M solution is to be diluted to 500 mL with a final concentration of 0.150 moles/L.
What volume of the stock solution is required?
Type your answer here
5. A stock solution of 10.0 M NaOH is prepared. From this solution you need to make 250.0
mL of 0.375 M solution. How many mL will be required?
Type your answer here
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Laboratory Techniques for Environmental Technicians
Practical: Metal colours in a flame
Background: Atoms when excited by flame, electric current etc are known to become
excited. When this occurs electrons jump to higher energy levels and on cooling fall back
into vacant electron holes. This results in energy, often in the form of light being emitted.
The colours emitted can be a useful identification tool.
Purpose: To observe the colours of various metal ions in solution
Procedure: (This will be done at the intensive session)
Aspirate each of the known metal ion solutions into the flame and note the colour of the
flame.
Aspirate the unknown metal solutions and identify from the colour of the flame the metal in
solution
Results:
Metal ion
Flame colour
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Laboratory Techniques for Environmental Technicians
Chapter 3: Safety in the Laboratory
Safety in the Laboratory
Purpose:
This section is designed to provide a brief overview of work safety in the laboratory. You will
need to refer to the text for laboratory safety for additional information.
Work health and safety laws have greatly improved safety and wellbeing in the workplace.
All employers must provide safe and healthy work conditions, all workers must work within
the safety systems and both groups must accept responsibility for identifying and controlling
hazards and minimising risk of harm. These laws are continually being reviewed and the
current Work Health and Safety Act 2011 resulted in a change from 1 January 2012.
Any laboratory has a range of hazards that need to be controlled to minimise the risk of
harm.
A working laboratory is really not much different to working in your home kitchen!
What is a hazard?
A hazard is something that can cause harm. For example a Lion is a hazard as it could kill a
person
What is risk?
Risk is the potential (likelihood) for a hazard to cause harm. For example, if your were a lion
tamer the risk of the lion would be high. However, if you were in a laboratory where there
are no lions the risk would be low.
What does hazard control mean?
This means that after identifying the hazard and the risk that appropriate control measures
are used to either eliminate the hazard or reduce the risk.
Hazard control uses a principle known as the Hierarchy or Control in order to control an
identified hazard
Explain the hierarchy of control
Type your answer here
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Laboratory Techniques for Environmental Technicians
Laboratory hazards
Physical
Chemical
Biological
Electrical
Explosives
Infective agents – bacteria,
viruses, etc
Mechanical
Corrosives
Animal and plant toxins
Manual handling
Flammable gases
Allergens
Cuts
Flammable liquids
Sensitisers
Burns
Flammable solids
Biological active substances
Radiation
Radioactivity
Answer
Compressed gases
Oxidisers
Answer
Mixes of the above
Answer
Dust
Mixes with other types of
hazards
Answer
Confined space
Poisons
Answer
Stress
Noise
Note that there are acute vs
chronic issues, dose and Answer
response issues, exposure
standards, monitoring /
Answer
awareness problems
Light
Answer
Of the above laboratory hazards are there any that also apply to your kitchen? Use an
asterisk * to indicate those that also apply to your home.
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Laboratory Techniques for Environmental Technicians
All chemicals should be considered to be dangerous because they may be:
◗ toxic / poisonous : damage body organs or tissue by interfering with the normal
chemical processes occurring in those tissues
◗ irritating: cause a bad response to eyes, skin and other organs
◗ flammable : burn or cause heat damage to tissue
◗ explosive:
create a shock wave of high speed and high pressure that causes
mechanical damage
◗ corrosive : chemically degrades tissue at the point of contact.
Specific organs such as lungs, liver, eyes and skin may be affected or gross damage can occur
at all points of contact. Damage can range from:
◗ slight tissue damage or irritation to total destruction
◗ minor, temporary change to widespread permanent damage
◗ immediate evidence of symptoms to long-term delay before onset.
◗ Effects may be acute (quick acting) or chronic ( build up over time).
A chemical is only dangerous if it enters your body. This can only occur by:
◗ inhalation
◗ absorption
◗ ingestion
◗ injection
What controls would exist for each of these chemical hazards in a laboratory?
Type your answer here
The law (Work Health & Safety 2011) now requires information sources about hazardous
chemicals to be supplied. They are called Safety Data sheets (SDS) and they help you to find
out about dangerous properties of chemicals and what measures to take to reduce harmful
effects.
Typically a SDS provides information about:
◗ common names and identification codes for the material
◗ physical properties
◗ major hazards of the material
◗ acute and chronic symptoms of exposure
◗ exposure standards
◗ medical advice
◗ spill and other emergency responses
◗ others
Search the ‘Web” for an SDS for concentrated hydrochloric acid and identify
information relating to the hazardous nature of the material.
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Laboratory Techniques for Environmental Technicians
Laboratory rules and regulations
◗ Always wash your hands before you leave the laboratory
◗ Only supervised work is permitted in the laboratory
◗ Chemicals and equipment are not to be taken from the laboratory
◗ Safety glasses, laboratory coats and appropriate shoes must be worn in the laboratory at
all times
◗ Clean-up spills immediately
◗ Act responsibly - the laboratory is not a playground, racetrack or amusement parlour
◗ Long hair must be tied back
◗ You must be aware of the location and operation of safety equipment
◗ All accidents and incidents must be reported
◗ Consult the SDS for unfamiliar chemicals
◗ Spillage of any chemical on the skin or eyes should be immediately treated with copious
quantities of water and the supervisor’s attention sought
◗ Eating and drinking in the laboratory is banned
◗ Fume cupboards should be used for work involving dangerous gases or vapours
The Laboratory is a safe place to work if you follow the rules
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Laboratory Techniques for Environmental Technicians
Practical 1.1 Laboratory Layout
Date completed: _____________________ Teacher Check _______________
Purpose
Analyst signature ____________
To become familiar with the layout of the work area, in particular, those areas and pieces of
equipment that deal with safety. Procedure:
Draw a map of the laboratory, which shows the location of the following features:
Fire extinguishers, fire exits, fire control equipment; special storage cupboards, safety
showers, eye wash stations, first-aid, fume cupboards, antidotes, laboratory store, ovens,
balances, waste disposal facilities.
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Laboratory Techniques for Environmental Technicians
Clean-up of spills
Common spills which can occur in the laboratory include solid reagents, solutions of all
kinds, various organic solvents and special ‘nasties’ such as mercury. All spills are to be
treated according to the immediate dangers. For example it may be necessary to:
◗ evacuate (e.g. poisonous gas or vapours)
◗ ventilate (e.g. organic solvent fumes)
◗ isolate services (e.g. flooding around electrical services)
◗ contain (e.g. prevent spread of fluid or other material beyond the site of the spill)
◗ absorb (e.g. use an inert absorbent to soak up a spill)
◗ neutralise (e.g. concentrated acids and bases need to be neutralised first to minimise
corrosion to the cleaning equipment and the cleaner)
◗ clean up (e.g. sweep up the solid residues)
◗ dispose (e.g. may need a clean-up service because of the hazardous wastes)
◗ rehabilitate (e.g. surfaces may have been damaged and will require refinishing).
The Golden rules for personnel safety in the laboratory:
◗ Leave the laboratory clean and tidy
◗ Put your equipment away
◗ Clean up all your mess
◗ Keep your equipment and community property clean and well maintained
◗ If you borrow anything, return it
◗ Any safety and hazardous situations or issues should be drawn to the supervisor’s
attention.
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Laboratory Techniques for Environmental Technicians
Practical 1.5 Safety in the laboratory
Date Completed:
_____________________ Teacher Check ______________
Procedure:
Analyst signature ____________
Observe the demonstrations on safety in the laboratory
Give a brief description of the demonstration
List the observations you made during the class and in the next laboratory session
Comment on aspects of the demonstration that you found interesting, disturbing,
informative etc
Results:
Test/
Demonstration
Observation
Hydrochloric acid
Session 1
Comments
HCl
+
Meat
Session 2
Sulfuric acid
Session 1
H2SO4
+
Meat
Session 2
Session 1
Nitric Acid
HNO3
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Laboratory Techniques for Environmental Technicians
+
Session 2
Meat
Session 1
Sodium
Hydroxide
NaOH
+
Session 2
Meat
Session 1
Ammonium
hydroxide
NH4OH
+
Session 2
Meat
Session 1
Propanone
Acetone
CH3COCH3
+
Meat
Session 2
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Laboratory Techniques for Environmental Technicians
Session 1
Session 2
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Laboratory Techniques for Environmental Technicians
Dangerous goods 'Class' labels
Class 1 - Explosives
Class 2 - Gases
Class 3 - Flammable liquids
Class 4 - Flammable solids
Class 5 - Oxidisers
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Laboratory Techniques for Environmental Technicians
Class 6 - Toxic and infectious substances
Class 7 - Radioactive material
Class 8 - Corrosives substances
Class 9 - Miscellaneous dangerous goods
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Laboratory Techniques for Environmental Technicians
Chapter 4: Basic Laboratory Equipment
Basic Laboratory Equipment
Laboratory testing is used to answer at least one of these basic questions:
◗ “what is this material?”
◗ “what are the components in this material?”
◗
“how much of each component is in this material?”
In other words, the lab can analyse a sample to describe what it is or what is in it –is there:
◗ mercury in this fish?
◗ salmonella in this chicken meat?
◗ cocaine in this powder?
◗ alcohol present on this person's breath?
The lab can analyse a sample to describe how much of any component is in it:
◗ how much mercury in this fish?
◗ how much salmonella in this chicken meat?
◗ how much cocaine in this powder?
◗ how much alcohol present on this person's breath?
The ‘what” questions are questions about quality and such laboratory work is called
qualitative analysis.
(if you answer 'yes' to the 'what' questions above, what does it tell you about quality of each
sample?).
Qualitative analysis requires qualitative equipment and the major focus is on isolating and
purifying components and identifying components.
Accuracy in measuring mass and volume is not a big deal because we are not answering
anything about ‘how much’. It is quite acceptable to use measuring cylinders and 2- and 3decimal place balances.
Accuracy in measuring physical properties, eg RI, pH, m. pt., etc is very important because
we use these numbers to decide on purity and identity.
The ‘how much” question is a quantity question and such laboratory work is called
quantitative analysis
Equipment for Handling chemicals must be chosen considering the purpose required:
◗ storage or
◗ transfer or
◗ processing or
◗ measurement
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Laboratory Techniques for Environmental Technicians
You will work with many types of equipment in the intensive session. The following is just
some of the equipment commonly found in a general ‘wet’ laboratory.
Storage containers
For large volumes:
◗ aspirators
◗ drums
◗ Winchester bottles (2.5L)
For smaller volumes and also for transfer:
◗ Conical flasks
◗ Beakers
◗ Conical beakers
Measuring glassware (both quantitative and qualitative):
◗ Measuring cylinders
◗ Pipettes
◗ Burettes
◗ Volumetric flasks
Miscellaneous equipment:
◗ Spatulas
◗ Funnels
◗ Evaporating basins
◗ Test tubes
◗ Mortar and pestle
◗ Tongs
◗ Balances
◗ Weigh boats
◗ Bunsen burners
◗ Heating mantles
◗ Desiccator
◗ Thermometers
◗ Retort stands
◗ Condensers
◗ Various hoses and tubing
Do you know any more to add to the list?
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Laboratory Techniques for Environmental Technicians
The following table lists basic apparatus used in the laboratory. Complete the table by giving
a use for the equipment and any safety features that need to be considered.
Diagram
Name
Use / Safety
Test tube
Measuring cylinder
Watchglass
Graduated and bulb
pipette
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Laboratory Techniques for Environmental Technicians
Diagram
Name
Use / Safety
Burette
Beaker
Conical flask
Evaporating basin
`
Crucible
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Laboratory Techniques for Environmental Technicians
Diagram
Name
Use / Safety
Pipe clay triangle
Tripod
Gauze mat
Volumetric flask
Filter funnel
Buchner funnel
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Laboratory Techniques for Environmental Technicians
Diagram
Name
Use / Safety
Vacuum flask
Spatula
Dropper
Tongs
Condenser
Mortar and pestle
Desiccator
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Laboratory Techniques for Environmental Technicians
Chapter 5: Introduction to Material Handling
In any laboratory, certain techniques are regularly performed and are essential to most
routine work. These techniques include the measurement and use of chemicals and other
hazardous substances and the operation of laboratory equipment. A competent technician
uses these basic techniques, confidently and without risk, in routine work and with any new
procedures. Procedures include:
◗ material transfer
◗ weighing
◗ volume measurement
◗ solution preparation
◗ filtration
◗ heating, cooling and drying
◗ temperature measurement
◗ using gas cylinders
◗ particle size reduction
◗ cleaning and drying glassware.
Material transfer
Materials are routinely moved from one container to another during laboratory work. The
transfer usually has to meet at least one of the following requirements:
◗ safety with materials which are toxic, corrosive or dangerous in some way.
◗ quantitative transfer where any loss will affect the accuracy of the task.
◗ contamination control either from the surroundings to the material or vice versa
Solid transfer
◗ usually handled by a spatula or weighing vessel or in the original container.
◗ If the solid is to be made up into solution then a solvent may be used to assist the
transfer
◗ The transfer must always be quantitative
Liquid transfer
◗ Generally liquids pour well
◗ A small amount of residue may be left in the container and depending on the use of the
liquid this may need to be washed into the next container
◗ Volumetric glassware has very specific requirements for transfer and these will be
discussed later.
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Laboratory Techniques for Environmental Technicians
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Laboratory Techniques for Environmental Technicians
Practical work 5.1 Basic material handling techniques
Date Completed: ______________________ Teacher check ________________
Purpose:
Analyst signature ____________
All chemicals whether solid, liquid or gas must be handled with caution. This practical looks
at the techniques to be employed when transferring solids and liquids.
Procedure:
Observe the demonstrations of material transfer shown by your trainer. Identify the
mistakes and indicate how the process should be completed.
Procedure
Errors
Corrections
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Laboratory Techniques for Environmental Technicians
Volume measurement techniques
Good laboratory practice aims to:
◗ control spills and dribbles
◗ avoid contact with hazardous contents
◗ reading the scales on the glassware from the meniscus location
◗ obtain quantitative measurement, transfer and delivery without loss or error
◗ make the correct choice of quantitative and qualitative glassware
◗ develop a sense of how valid a transfer or a measurement has been.
Liquid Transfer Techniques
Transferring the liquid to a new container (receiver) needs to avoid:
◗ spillages (missing the receiver)
◗ dribbles (some liquid runs down the outside of the original container)
◗ losses (not all of the liquid reaches the receiver).
Reading the Meniscus
Measuring liquid volume requires you to read the position of the meniscus on a scale. The
meniscus is the curvature or shape adopted by a liquid surface near the walls of any
container. It can curve up or down. For example, mercury in glass, is concave down whereas
water in glass, is concave up. All volume judgments are made by comparing the flattest
portion of the meniscus with the scale on the vessel. With water this becomes the bottom
of the meniscus and with mercury it is the top. The volume scale may be a single mark
(pipette, volumetric flask) or a continuous scale (burette, graduated pipette, measuring
cylinder) which can cause trouble if the meniscus lies between two markings. You should
only guess to half of the smallest division on the scale.
Accurate and Approximate Volume Measurement
When accuracy in volume measurement is the first concern, then volumetric glassware
(quantitative glassware) must be used with the correct technique and the right attitude. The
professionalism which tells an analyst whether a volume measurement has been performed
validly or not, comes from:
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Laboratory Techniques for Environmental Technicians
◗ proper understanding
◗ technical skills
◗ the correct attitude.
In particular an attitude of near enough is good enough is not acceptable. Sometimes the
analyst knows the volume measurement or transfer has been faulty and must make the
professional decision to discard the faulty material and repeat the necessary steps. The
word aliquot is commonly used to denote an accurately measured volume of liquid and will
be used frequently in this text in the section on titration.
For those volume measurements where the amount needed is not critical, the following
may be appropriate or convenient:
◗ a qualitative glassware such as measuring cylinders
◗ graduated beakers or flasks
◗ any of a wide range of droppers fitted with a suction bulb (e.g. Pasteur pipettes,
droppers fitted to reagent bottles and even droppers fitted with graduated stems)
◗ an experienced guess.
Nevertheless, you should still pay attention to safety and proper transfer technique and
avoid parallax errors.
Types of droppers
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Laboratory Techniques for Environmental Technicians
Pipetting Techniques
All single-mark pipettes are used by filling them to an upper mark. A sacrificial rinse (three
times) with the fill solution is the minimum effort all users must make to ensure the pipette
is clean. This must be done even if the pipette is dry or after normal cleaning. The sacrificial
rinse only removes fresh contaminants or solvents which may have entered during current
usage. The entire contents are then delivered to the receiving vessel
General pipette rules
Perform sacrificial rinse with the fill solution.
Overfill to above the mark.
Set the bottom of the meniscus on the line.
Allow to drain in a vertical position until no more liquid runs out.
Touch the tip on the side of the receiver to draw the final amount of liquid out.
Do not blow it out. There must always be a little left behind, which is allowed for in the
original calibration.
Overfilled
Too low
Meniscus is
spot-on
The pipette is
not yet empty
Too high
The pipette has
been properly
emptied
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Laboratory Techniques for Environmental Technicians
Burette Techniques
All burettes have a highly graduated scale which starts from 0.00 mL at the top and goes
down to the biggest volume at the bottom. A sacrificial rinse with the fill solution is the
minimum effort all users must make to ensure the burette is clean. This must be done even
if the burette is dry or after normal cleaning. The sacrificial rinse only removes fresh
contaminants or solvents which may have entered during current usage. It is normal to use
a funnel to fill burettes but the funnel must be removed before zeroing the burette as liquid
is often trapped between the funnel and the burette and tends to escape half-way through
a job and ruin the measurement in progress.
Burettes can have a range of taps at the bottom, from the traditional glass stopcock (which
needs a touch of lubricant and care not to pull it out as it is being used), to more modern
teflon stopcocks and rotoflow-type screw taps which are much easier to use, do not come
out easily and are self-lubricating. Techniques for using burettes are shown in Figure 5.5.
General burette rules
◗ Perform sacrificial rinse with the fill solution.
◗ Overfill to above the mark (zero or any other).
◗ Open the tap fully and ensure all air bubbles are driven out of the tap and the tip of the
burette.
◗ Set the bottom of the meniscus on the chosen mark — usually zero.
◗ Clamp in a vertical position at a convenient height above the receiver.
◗ Touch a wastes receiver onto the side of the burette tip to draw any excess liquid off the
tip.
◗ Check your mark and record its value.
◗ Run the required volume of liquid out and ensure any suspended drop on the burette tip
is also transferred.
◗ Read and record the new volume. Perform any necessary subtractions.
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Laboratory Techniques for Environmental Technicians
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Laboratory Techniques for Environmental Technicians
Volumetric Flask Techniques
All volumetric flasks have a single mark and hence contain but a single known volume and
all are used by filling them to this mark. A sacrificial rinse with the intended solvent is the
minimum effort all users must make to ensure the volumetric flask is clean. This must be
done even if the flask is dry after normal cleaning. The sacrificial rinse only removes fresh
contaminants or solvents which may have entered during current usage.
The primary purpose of volumetric flasks is to make up solutions or dilutions of solutions.
General volumetric flask rules
◗ Perform sacrificial rinse using intended solvent.
◗ Add the solute first, quantitatively, without loss. For many solutes, the transfer into a
small beaker for dissolving in the solvent is a useful technique to avoid loss of solute.
The solute may be solid or liquid or a solution already.
◗ Add some solvent to ensure the solute dissolves completely.
◗ Transfer carefully with pouring and rinsing to the volumetric flask
◗ Mix well by inversion and swirling.
◗ Solvent is added until the bottom of the meniscus is on the line — not over, not under —
but on the line. Mix well and support the weight correctly.
◗ Label with name, concentration, date and appropriate safety advice.
You cannot remove any liquid if you overshoot the line when making up to volume. You will
also remove an unknown amount of solute and hence invalidate the job.
The last step requires a final mixing which is done by inversion and swirling, using the
trapped air bubble to ensure mixing occurs. The lid and flask bottom must be solidly
supported or a spill or a breakage will result.
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Laboratory Techniques for Environmental Technicians
When the lid is removed to access the contents, you must take care to avoid contamination
of the contents by the laboratory environment and vice versa from a carelessly placed lid.
Measurement Of Mass
Mass in the everyday world is measured on scales such as those used for people or
vegetables or fertiliser. Bigger versions are used to check trucks or bales of wool and smaller
versions are used by the jewellery trade. In the scientific world, the term ‘scales’ is replaced
by the term ‘balance’ when describing tools for the measurement of mass.
Laboratory Weighing
Weighing of materials is fundamental to most procedures. You should be competent at
working with each of the balances found in the general laboratory.
Weighing rules
◗ all balances must be clean, level and zeroed
◗ all chemicals must be weighed by difference in a container
◗ don’t weigh hot or wet objects
◗ clean up all spills immediately
◗ check the balance has returned to zero and reset it before you leave.
Common terms used with weighing or the use of a balance include:
◗ clean: nothing on the balance pan
◗ level: the workbench and the balance need to be horizontal
◗ balance reading: the mass on display
◗ true mass: the mass of the sample after the container mass has been subtracted
◗ reset: leave the balance clean, level and zeroed
◗ weighing by difference: an empty container is weighed, the sample is added to it and
the mass of the container and material recorded, the difference between the final mass
and the empty container is the mass of the material
◗ taring: the balance pan is adjusted to read zero with the empty container sitting on the
balance pan, when the material to be weighed is added, the displayed weight is the
weight of the added material only
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Laboratory Techniques for Environmental Technicians
Practical work 5.2A Introductory weighing task
Date Completed: ______________________ Teacher check ________________
Purpose
Analyst signature ____________
The aim of the practical is to become familiar with the different types of balances available
in a general laboratory. Your instructor will demonstrate the correct procedures to follow
when using a balance.
Procedure:
Use the range of laboratory balances supplied as shown during your teacher’s
demonstrations. You will perform repeated measurements which will be used to check your
accuracy and that of the balances. Record the capacity and sensitivity of the balances.
Results:
Balance
Object Code ........
Object Code........
Object Code ........
Triple Beam:
Capacity: ...........
Sensitivity;..........
3 dp top pan
Capacity: ...........
Sensitivity ..........
4 dp analytical
Capacity ............
Sensitivity .........
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Laboratory Techniques for Environmental Technicians
Other Basic Laboratory Procedures
Separation of one component from another in a mixture can be a routinely necessary task in
a laboratory. Filtration is used to separate an insoluble solid from a liquid. The procedure
requires the use of a barrier called the filter, which is permeable to the liquid but retains the
solid. The liquid portion is termed the filtrate or mother liquor and the solid is the filter cake
or residue.
Two common approaches to filtration are:
◗ gravity
◗ vacuum assisted
Gravity filtration is a simple method of separating a solid from a liquid. It involves the use of
a filter paper, available in various pore size and diameter, folded and placed in a filter
funnel. The mixture is poured in the top, the paper allows the liquid to pass and the residue
is left behind.
Vacuum filtration is used where large quantities of solids and/or fluids are involved and fast
separation is required. Problems arise if the solid is very fine or gelatinous.
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Practical work 5.5A Simple filtration
Date Completed: ______________________ Teacher check ________________
Purpose: To become familiar with simple filtration techniques, following instructions and
noting safety requirements.
Analyst signature ____________
Note: The chemicals in this practical are very toxic. Ensure that any spills are reported and
also ensure that hands are properly washed if skin contact occurs and before leaving the
laboratory.
Record the identity and quantity of each of the chemicals
Dispense an aliquot for each reagent bottle into a clean dry beaker
Carefully mix, using a glass stirring rod, the beaker contents over a steam bath for five
minutes
Cool the beaker in an ice bath
Set up a filtration apparatus as shown by your teacher and using a labelled, pre-weighed
filter paper filter your sample.
Wash the filter cake well with distilled water
Transfer the filter cake and filter paper to a labelled watch-glass and allow to dry
Reweigh and determine the amount of solid material.
Results:
Identity of solution 1
Volume of solution 1 used
Identify of solution 2
Volume of solution 2 used
Mass of empty filter paper
Mass of filter paper + precipitate
Mass of precipitate
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Questions:
When is filtration a suitable method of separation?
One of the solutions you worked with contained lead. Why is it necessary to be extremely
careful when working with this chemical? How would you find out the necessary safety
requirements for working with lead solutions?
Why was it necessary to stir the solution over a water bath?
Why was it necessary to allow the solution to cool before filtration?
When filtering the solution is not filled to the top of the filter paper, why?
Draw a diagram to represent the simple filtration apparatus.
List all the possible places where losses may have occurred.
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Laboratory Techniques for Environmental Technicians
Heating in the laboratory
A laboratory also will have an array of heating, cooling and drying equipment such as gas
burners, heating mantles, muffle furnaces, hot plates, steam baths, ovens, desiccators and
refrigerators.
Common gas burners include the bunsen, the meker, batswing and microburner. These
burners all depend on a supply of fuel gas and air. The efficiency and heat output of any
burner can be controlled by selecting the amount of air and/or gas which passes through
the burner
.
Bunsen burner
Batswing burner
Meker burner
Your teacher will demonstrate the correct ignition procedure for the common laboratory
burners . Generally if you are going to leave a bunsen burner alight, but not in use, the
yellow “safety” flame should be burning.
Non-flame devices are versatile and much safer for use with flammable materials. These
include electric hotplates, drying ovens, muffle furnaces, heating mantles and drying
cabinets.
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Laboratory Techniques for Environmental Technicians
Practical: 4.4 Heating devices
Date Completed: ______________________ Teacher check ________________
Procedure:
Analyst signature ____________
Examine the heating equipment on display and complete the following table.
Device
Description
Major Use
Safety Hazards
Flame dependent
devices
(Burners)
Bunsen
Meker
Batswing
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Laboratory Techniques for Environmental Technicians
Microburner
non
–
devices
flame
Heating mantle
Electric hot plate
Muffle furnace
Laboratory drying
oven
Heat lamp
Steam / water /
sand baths
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Laboratory Techniques for Environmental Technicians
Others
Practical:
Application of heating equipment
Date Completed: ______________________ Teacher check ________________
Procedure: Follow the instructions for each task and record the data where required.
Analyst signature ____________
Task 1
Unscrew the stack on each of the supplied burners and examine the features of each
burner. Note the gas jet or nipple and the rotatable sleeve which admits air and mixes it
with the gas as they moves up the stack.
Attempt to light the gas at the nipple. Note - the gas supply may need to be very low for
this.
Replace the stack and light the Bunsen as demonstrated by your teacher.
Note the effect of sleeve positioning on the appearance of the flame.
Using the blue flame, turn the gas pressure down slowly so that the flame gets smaller and
strikes back down the stack to burn at the gas nipple.
This condition is extremely dangerous because:
the Bunsen appears to be off and hence is a fire hazard
the stack will be heated by the flame and will burn anyone who touches it
6. Repeat the above five steps with the Meker and micro-burners.
Draw and label a Bunsen and discuss your observations below
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Laboratory Techniques for Environmental Technicians
Task 2
Use the supplied thermocouple to measure the temperature of various zones in both the
‘blue’ and ‘yellow’ versions of the Bunsen flame.
Complete the following table:
Zone of the flame Yellow Flame Blue
Flame Diagram to locate each zone
being measured
temperature
temperature
(OC)
(OC)
Approx.
20cm
above the tip of
the flame
At the tip of the
flame
20
cm
In the heart of the
flame
(top of blue cone)
Level with
stack opening
the
Task 3
This section is designed to investigate the changes which occur to the gas mixture in the
flame.
Following the demonstrations given by your teacher, examine the presence of zones of
unburnt gases in a Bunsen burner as follows:
Use a thin glass tube inserted directly into the bottom of the blue cone to tap off some of its
contents – attempt to ignite the gases at the other end of the tube
Pierce a live match with a pin inserted at right angles approximately half way along the
match – suspend this match with its head pointing up, on the stack of an unlit Bunsen. Light
the Bunsen, with the air hole open.
Observations
Record your observations and provide an explanation of what is happening.
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Laboratory Techniques for Environmental Technicians
Chapter 6 The dreaded calculations
This chapter has an online attachment that you may find useful for the calculations. The
assessment for the unit is composed of 2 open book assignments that are located in the
appendix of this theory booklet.
Laboratory Data and Error Analysis
How can you know if you have the right answer? Unfortunately, there is often no simple
way to know whether the correct answer has been obtained. To improve our confidence,
scientists have developed ways to estimate the uncertainty in the measured value(s) and
how they differ from the true or correct answer(s). Laboratory scientists use a series of
terms when describing uncertainty and errors.
Accuracy
A measurement is accurate if the correct answer was obtained. The correctness of an
answer depends on what standards have been set (i.e. how close to the true answer is
acceptable). Hence every measurement will have some degree of accuracy and also a level
of error. Scientists also talk of accurate instruments or inaccurate methods because of the
quality of the instrument’s design or the limitations of the method.
Error
The amount by which the true value has been missed is called the error. Some errors are
brought about by the limitations of the equipment. Analysis of these limitations can be used
to calculate the size of the errors. Other errors are linked to limitations in the skill or
vigilance of the operator or shortcomings in the method. These errors are much more
difficult to identify and compensate for.
Precision
If very similar answers (a standard for closeness needs to be set) are obtained for a set of
measurements (e.g. duplicate, triplicate, replicate determinations), the answers have a high
level of precision. Precision can refer to the results of a person or a testing method or a
laboratory. Measurements which lead to answers which are highly precise, are often taken
as evidence that the answers are also accurate. However if the same error was committed
each time, good precision may still be all wrong.
Reliability
Repeated performance of a procedure which achieves good results every time, means it is
reliable and its errors are under control. This is used to describe the quality of a technician
or a measurement or a technique or a method. Reliability is also known as repeatability.
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Laboratory Techniques for Environmental Technicians
Validity
A method or procedure must be able to be supported by traditional scientific reasoning as
well as the good data it produces. For example, a number of alternative healthcare practices
exist which are labelled as invalid because traditional practitioners can’t see a scientific link
between the alternative approach and current understanding of medicine. The
improvements in the patient can’t be taken as proof of the validity of the alternative
approach.
Classification of Errors
All scientific work is subject to errors and the skill of a competent worker is to understand
where errors may come from and how to minimise or compensate for them. Errors normally
can only arise from defects in:
◗ the operator
◗ the equipment
◗ the method being followed.
Operator errors are those errors for which the operator is responsible and can be caused by
physical handicaps (e.g. colour blindness), bias, prejudice or poor attitude to quality and
accuracy. Examples of operator errors include:
◗ mechanical loss or gain of materials during weighing, filtering, liquid transfer
◗ failure to obey essential conditions such as temperatures, times or conditions
◗ incorrect performance of required techniques such as dilution, titration, instrument
optimisation and operation.
Equipment errors are due to defects in the analyst’s tools and equipment or the effects of
environmental factors upon them. Common equipment errors are:
◗ random defects such as a balance or oven not operating as expected
◗ systematic defects producing results which appear acceptable but are consistently in
error; they are caused by the use of tools such as weights, graduated glassware and
thermometers which are out of calibration
◗ equipment which is not appropriate for the task; you needed to use a different size or
model or system
◗ reaction of reagents and samples with glassware and other containers, resulting in the
presence of interfering materials
◗ use of reagents containing impurities.
Method errors are due to errors in the procedure or the technique. Common method errors
are:
inaccuracy — the method used is not adequate to get the measurement needed;
measurement of pH using indicators or test strips may be inaccurate because the conditions
in your sample interfere with the performance of the indicators
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Laboratory Techniques for Environmental Technicians
lack of validity — the method cannot cope with limitations of your sample; moisture in
many fresh foods cannot be validly obtained by simple oven drying; lead in blood or
pesticides in meat require extremely sensitive testing equipment and cannot be done by
simple test-tube reactions
calculations and processing of data which have been incorrectly carried out
incorrect sampling or sample preparation
failure of the necessary reactions to go to completion
occurrence of side reactions and by-products.
Another system uses the terms systematic (or determinate) errors and random (or
indeterminate) errors to classify the errors which affect an experimental result.
Systematic errors are those that have a definite value which can be measured and
accounted for, that is, they are errors that are possible to avoid, minimise or compensate
for.
Random errors result from the person, the equipment or the method operating outside its
limitations. These errors may or may not be positively identified and so will not always have
a definite measurable value. Random errors occur in a random manner.
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Laboratory Techniques for Environmental Technicians
Systematic and random errors can be minimised by:
◗ calibration of apparatus and application of correction factors
◗ use of blanks, spiked samples, controls, certified reference materials, recovery tests and
variations in sample size
◗ alternate or independent methods of analysis
◗ replicate determinations
◗ standard addition techniques
◗ internal standard techniques
◗ validation using the literature or prior work with similar samples
Basic Numerical Methods for Reporting Measurements
When any scientific measurement is carried out, it is not enough to simply carry out one
analysis, and to report the result. There are too many sources of possible error that can
arise, and therefore the analysis should be repeated a number of times to confirm that the
measurement obtained is reliable. The result that is reported after several, acceptable,
measurements of the same property is the average of the repeats or the mean. This is
calculated from the sum of the answers divided by the number of values. For example, an
analyst determines the vitamin C content of a sample and she obtains the following results:
24.39 mg/L,
24.18 mg/L,
24.27 mg/L.
The reported value for this analysis would be the mean of these determinations, that is:
( 24.39 + 24.18 + 24.27 ) / 3 = 24.28 mg/L.
The two attributes needed to accept these data are:
accuracy (i.e. the result should be as close as possible to the actual value)
precision (i.e. the replicate values should be close to each other).
In this case, if the accuracy could not be estimated, the precision would be considered good
and the analyst should feel confident with her work. When accuracy estimates are essential
(e.g. nutritional labelling, medical analysis), then a number of validation methods exist.
Routine validation methods which could be used are:
control samples — samples for which the amount of analyte (vitamin C) is already known;
the analysis for these must be correct to be able to accept the others
spiked samples — samples which have already been analysed and to which a known amount
of analyte (vitamin C) has been added; the increased level of analyte allows a new expected
result to be calculated and again the analysis for these spiked samples must be correct to be
able to accept the others
alternative methods or procedures — if the same answer is still obtained then both
approaches would be valid.
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Laboratory Techniques for Environmental Technicians
Relative and Absolute Descriptions of Error
The absolute error (or absolute accuracy) is the difference between the observed value and
the true value.
The relative error (or relative accuracy) is the absolute error expressed as a percentage of
the accepted value. The sign of the error may be positive or negative, indicating that
the result is high or low respectively.
The absolute precision is half of the range of the measurements.
The relative precision is the absolute precision expressed as a percentage of the mean of the
measurements.
These definitions are summarised in the Table below
Definitions and formulae used to describe errors
Symbol
Interpretation
Formula
X
A measured or observed value
x1 , x2, x3, etc. for all your readings
R
The range from biggest to smallest of all
replicates for this measurement
R = x biggest – x smallest
µ
The average of all replicates for this
measurement
µ = [x1 + x2 + x3 + … ]
number of replicates
X true
The true or correct value
Eabs
The absolute error or accuracy
Eabs = X – Xtrue
or
Eabs = µ – Xtrue
Erel
The relative error or accuracy
Erel =
Eabs × 100
Xtrue
Pabs
The absolute precision
Pabs =
Prel
The relative precision
Prel
R  2
= Pabs × 100
µ
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Errors and Limits of Reading
The uncertainty of any reading is half the value of the smallest scale division. Thus if your
smallest scale division is one unit your answer can be shown as  0.5 unit.
Reading scales and assigning uncertainty
Read the scales below and write down your readings for the position of the pointer in each
case. There are no units required. Assign an uncertainty figure to each of your scale
readings. Use  after the value for each reading.
a
b
h
n
o
c
d e
i
p
j
q
k
r
f
l
g
m
s
t and u
(a)
(h)
(o)
(b)
(i)
(p)
(c)
(j)
(q)
(d)
(k)
(r)
(e)
(l)
(s)
(f)
(m
)(n)
(t)
(g)
(u)
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Laboratory Techniques for Environmental Technicians
Chapter 7: The why’s of sampling
Introduction
Sampling is a process where portions of a material are taken for testing. The test results are
then assumed to apply to the entire material.
Scientific testing absolutely depends on linking the laboratory results for a sample to the
bulk of the material from which it came. If there is no valid link, there is no point in the test.
How would you, as a science technician, advise authorities about a testing system for the
analysis of drinking water from a river or dam? Three different approaches or options would
cover your choices:
Option 1 — test the water while it is still in the dam
Option 2 — take some water out and test it with field equipment
Option 3 — take some water out and send it to a laboratory.
Option 1 seems good because the water is undoubtedly representative and it is hard to see
it being changed by the test process. For example a pH probe or a dissolved oxygen sensor
can be dipped directly into the water from a boat and readings taken from as many different
sites as is needed.
Option 2 could be necessary when some sample treatment like heating, filtering and other
chemical additions are necessary. But the danger exists that the sampling will take a bad
batch of water or contaminate the water which is taken and so the test results may not
validly apply to the rest of the dam. But what other choice do you have?
Option 3 would be needed when field equipment can’t do the job and the controlled
environment of a laboratory is needed. The problems now include the time and conditions
which may cause the sample to undergo significant physical, chemical and biological
changes. Oils and greases may be lost on the walls and lid of the container, bacteria may
multiply or die, fine sediments may settle and absorb heavy metals or pesticides. Sounds
risky?
In reality, all approaches are used and all are satisfactory so long as the limitations are
understood and sampling is carried out properly. The best equipment, technical expertise
and hard work cannot compensate for a poor sample.
Sampling Requirements
Some materials are very easy to sample without bias because they are very uniform. Others
are more complex which makes sampling very difficult to perform correctly.
Liquids such as:
◗ drinking water
◗ milk
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◗ blood
◗ diesel fuel
would be tested using a sample which was only a tiny fraction of the whole, and the results
would be very confidently applied to the entire bulk of the material. Uniform or
homogeneous materials are usually free of bias in sampling (but are still liable to problems
with dirty containers and poor storage conditions).
When all portions of a bulk material are not identical (the material is heterogenous), it
becomes possible to collect biased samples. The earlier liquid examples could become:
◗ muddy water
◗ curdled milk
◗ clotted blood
◗ diesel fuel floating on a layer of water.
and so the sampler needs to know whether to skim some good stuff off the top or just grab
a bucketful and hope for the best. Obviously the law does not allow us to just hope for the
best when human health and safety are at risk and a better system is needed. In general,
the better system has been developed after much trial and error.
Gases are usually well mixed (e.g. atmospheric and flue gas samples which are moving) but
layers can develop in confined spaces such as tanks and mines, where different gas
composition may build up on the floor or ceiling. Again, the sampling techniques will
depend on the uniformity expected in the bulk material. Experience, technical skills and
statistical competency are all needed to get it right.
Solid materials of natural origin (e.g. coal, agricultural produce and minerals) may or may
not be uniform and synthetic products such as building materials, processed foods, cosmetic
preparations and cleaning agents can have many different components blended into a
number of different phases and at varying concentrations. Sampling of these for testing in
order to meet the legal and quality and economic standards of the organisation is of great
importance to future survival and prosperity.
A final example to introduce the needs and difficulties of sampling is taken from an area of
global concern — our environment. The build-up of harmful contaminants and of dangerous
toxins in our air, water and soil is well known. Responsibility and control is difficult to
enforce because it is hard to prove:
how much pollution existed before (background levels of carbon dioxide, ozone,
radioactivity, heavy metals and coal and petroleum by-products)
who released the contaminants, since they quickly become dispersed or diluted or mixed in
with those from next door. Unless you catch them, it is hard to allocate blame correctly.
Sampling of contaminated and uncontaminated sites is a necessary step to answering such
questions, but how to perform the process to clearly identify the source is not easy.
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Sampling Terminology
A sample is a small portion of a large mass of material and must be representative of that
mass. The mass could be hundreds of tonnes of ore, coal or grain; thousands of litres of
liquid such as milk or petrol, or millions of cubic metres of natural gas, or even the whole
atmosphere. Quite clearly, analysis of the small portion should give an answer for the
concentration of each analyte to accurately reflect its concentration in the large mass which
it represents.
A representative sample must be identical in its chemical and physical characteristics to the
whole. Thus it must be taken correctly and stored and processed for testing in such a way
that its representative status is not affected.
A specimen refers to a portion or single part of the whole. It may or may not be
representative of the whole. Hence a nice specimen of coal can be picked up from a
stockpile because of appearance or size or convenience, but this alone does not guarantee
that it is a representative sample of this coal stockpile.
Homogeneity refers to the degree of uniformity in the composition and distribution of each
analyte in the bulk material.
Homogeneous materials are entirely uniform and any specimen is also a representative
sample.
Heterogeneous materials lack uniformity in their overall bulk in at least one property
(colour, particle size, chemical composition, hardness, crystallinity, etc.) and a specimen
may not be representative.
Well-mixed gases and liquids are considered to be homogeneous. Conditions are possible
however, where gases and liquids can form layers of different composition due to:
vapours and evaporation (e.g. petrol)
phase separation (e.g. silt in river water)
temperature (e.g. cooking oils)
corrosion effects in containers (e.g. confined spaces and oxygen depletion).
Solids can be homogeneous or heterogeneous depending on the component of interest.
Cement looks to be a uniform grey powder to the casual observer, but particle size
distribution is critical to its performance and it is considered to be heterogeneous. It
requires a rigorous sampling scheme to detect particle size violations. Potatoes on the other
hand look very heterogeneous yet potato chip manufacturers would not care too much
about size, shape or colour so long as they all contained very similar amounts of reducing
sugars to ensure they cook to the same final golden colour at the end. Coal is particularly
complex with a range of classification schemes linked to the end use (steam generation or
coke production). Strict specifications on the range of physical and chemical properties are
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set by buyers and ultimately these specifications dictate the price of the coal. Proper
sampling and testing of coal is critical to the commercial success of any coal supplier.
Random sampling suggests that the sampler should gather material in a widely distributed
pattern but the pattern should not bias the removal to only one particular type of material
(e.g. the loose bits on the floor that would otherwise need to be swept up). The design of
the sampling pattern is critical to the confidence levels of the overall findings. Each industry
needs to carefully design its random sampling scheme and samplers need to understand the
effects of departure from this scheme.
Bulk (or gross) sample describes the end result of the collection of material in a sampling
program. The size of this sample can be quite large (hundreds of kilograms) and often needs
to be subsampled to obtain laboratory samples.
Subsampling refers to the process used to reduce the size of a sample in a representative
manner so as to obtain a more convenient quantity for laboratory work or storage. A variety
of manual (cone and quarter) and mechanical (riffle, sample splitter) methods are available
to ensure that this is done representatively. These are described later on.
A laboratory sample is the portion of the bulk sample provided to the laboratory for its
testing purposes. It has strict labelling requirements to ensure it can be linked to the original
supply. It must be preserved to ensure its properties are not changed by storage or
handling.
An analytical sample is that portion of the laboratory sample which is actually tested. It still
needs to be representative and also may have pretreatment requirements in terms of its
physical and chemical state.
Sample preparation is a description of the later stages of sample handling. Typically this
may require grinding or milling, sieving, drying, filtering, cooling and a variety of other
chemical and physical treatments. The figure below shows an idealised relationship
between your bulk supply and the sample tested by the laboratory.
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Samples taken to represent the bulk supply
A bulk supply of material can
have the following variations:







identity
size
state
uniformity
danger
accessibility
value
A laboratory sample of
adequate size for
testing and back-up
Analytical samples
which are labelled
and preserved for
testing and
storage
Relationship of
bulk supply to laboratory sample
Steps in sampling
Collect the gross or bulk sample from the material stockpile. Homogeneity of solid bulk
samples can be improved by coning (see Figure 6.3).
Reduction of the gross sample to a convenient size for laboratory handling. This is done by
coning and quartering, rifling, tabling, sample splitting, etc. (see Section 6.5).
Preparation of sample for analysis.
Repeated coning to homogenise a bulk solid sample
Obtain your individual samples
form the composite into a cone
Combine individual
samples into the
composite sample
Take away the first cone
by the edges and make
a new cone
continue until you form a
new cone and repeat the
coning process until you
have a satisfactory
homogenate
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The sampling process can present a number of dangers. There may be:
◗ physical dangers from the location (height, temperature, confined space)
◗ dangerous behaviour of the bulk material (unstable stockpile, dusts, volatility)
◗ hazards because of its chemical properties.
The sampler must wear adequate protective clothing, have a knowledge of the material
being sampled, and follow standard procedures laid down for the sampling programme.
◗ Sampling Equipment and Techniques
◗ Solids in general require a sequence of steps to be followed:
◗ collect the gross sample by random sampling,
◗ sub-sample the gross sample to create the laboratory sample,
◗ prepare the analytical sample by appropriate pre-treatment before analysis.
Large samples can be manually reduced into smaller representative portions by riffles,
coning and quartering, rolling and quartering, and other forms of sample splitting. Some
solids such as metals need to be drilled to obtain suitable samples and soil sampling in the
field may need core samples to be taken. Surface samples such as chips, clods and shavings
may not be representative.
A series of
chutes or
slides to split
and deliver
the gross
sample to a
A riffle box
set of new
locations or
collection
points
Gross sample is
poured into the chute
system at the top
Other chute patterns
Sub samples emerge and
are collected at the bottom
of each chute.
Riffle designs used to sub-sample the gross sample
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You have a laboratory sample of a suitable size
Or using a technique of coning and quartering to reduce the sample size
Figure 6.5
The quartering process used to subsample the gross sample
Liquids which are static and homogeneous are very easy to sample as any small portion will
represent the whole. A sample may be obtained by from a single point (called a grab
sample) using a container made from material appropriate for storage of that liquid. Liquids
with concentration gradients, separate phases and suspended or filtrable solids need more
thought, planning and equipment. These are normally sampled using devices such as diptubes, depth sampling bottles and sample thieves (see Figures 6.2 and 6.6). Piped liquids
which are sampled through outlet valves, can be extremely dangerous because of the
pressure and temperature of the emerging fluid.
Gases which may be sampled could be:
◗ toxic, explosive or flammable
◗ hot or cold
◗ high or low pressure
◗ contaminated with moisture or particulates
◗ unstable and require special preservation (eg. environmental air pollutants).
Many can be collected under their own pressure, but some sampling situations may require
a pump. Some gases are deemed to be of such concern that they are analysed by
continuous on-line equipment which automatically collects its own sample. For example, air
in underground mines can be explosive or poisonous or depleted in oxygen and needs
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constant careful monitoring. This equipment needs to be calibrated periodically for the
quantity of material taken and the response of the sensor system. Portable gas testers such
as stain tube analysers (also called Draeger™ tubes) and the explosimeter for flammable
gases are excellent field instruments, but some gases need to be collected, often over an
extended time period, by an absorbing medium (charcoal or reactive liquid medium) for
later analysis.
Sample Preparation, Preservation and Labelling
Some samples, even if taken in a valid and representative way, may still not lead to an
accurate analysis of the bulk material because of a number of difficulties. For example:
The sample may change or no longer represent the original, because conditions in the
sample are no longer the same as the bulk material. Examples include::
food taken from a hot or cold storage area (and so moisture levels, bacterial growth, vitamin
levels may be different),
a water sample from a pool or river or dam (a different temperature, dissolved oxygen and
light level will affect some components),
a blood sample (taken for blood alcohol or other component remains static but the person’s
biochemistry continues to function) or
a soil coring (drying out, contact with air leads to chemical and physical degradation).
The sample may have a chemical or biological composition which protects an important
component from being detected accurately. Silicates are notoriously insoluble and samples
with silica content need to undergo severe grinding and digestion procedures to ensure
their components are released for analysis.
The sample has chemical or biological components which are particularly sensitive and need
to be protected in some way.
The sample may need to be stored for a long period as a backup or because of delays in
analysis.
The only answer to all these difficulties is to apply a scientific approach to the perceived
problems. Each individual case needs to be assessed for where it may be breaking down, if it
can be fixed by reasonable means, and if you are able to report a valid upper or lower limit
for the analysis. Standard methods exist to address sample preservation and sample
preparation.
Sample preservation or protection methods include:
◗ cooling eg. food and biological samples
◗ acidification eg. heavy metals in water samples
◗ thiosufate additions eg. destroys chlorine to protect micro-organisms in water samples
◗ sterile containers and aseptic handling techniques
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Sample preparation is the procedure followed in treating the sampled material so it is
suitable for testing in the laboratory. The following steps may occur:
grinding, milling or blending as a way of reducing particle size and ensuring the necessary
components are released from the material
digestion where the sample is treated with acids or buffers or other solutions to ensure the
release of desired components in their correct chemical forms for analysis.
Labelling is the final requirement of the sampling program and needs to identify the
following essentials:
◗ date
◗ identity of the sampler and bulk sample
◗ the destination of the sample
◗ its testing requirements
◗ the destination of the analytical results.
Often there are a series of codes for these. Computerised systems are becoming routine for
reporting results. They ensure that sample results are properly logged, checked for
accuracy, and accounts and reports produced and despatched. Sometimes a reserve sample
is put into storage in the event of a dispute. Its labelling must allow it to be identified as part
of the original sample.
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Practical:
Date Completed:
Validation of sampling
___________________ Teacher check _____________
Purpose
Analyst signature ___________
This practical task is designed to demonstrate whether the sampling procedure used has
produced a sample that truly reflects the composition of the bulk material. You will create a
bulk supply of material of known composition (~5% salt in sand) and you will sample it for
laboratory testing. You will measure its true salt content and compare your answers to the
expected values.
Procedure
Mix sand (500 g) and sodium chloride (25 g) together to make a homogeneous blend by
coning on a sheet of plastic. Record the masses used in your logbook.
Sample your mixture by the cone and quarter technique until you have a laboratory sample
of approximately 25 g for testing.
Weigh three 5 g samples into 100 mL beakers for analysis. Record the actual masses in the
table below.
Analysis of samples
To each sample in a beaker, add distilled water (50 mL), mix well to dissolve the salt and
filter through a preweighed filter paper into a preweighed evaporating basin.
For each sample, use another portion of water (20 mL) to rinse the remaining salt residues
through the sand. Allow both washings to combine.
Evaporate the liquid in the evaporating basin on a steam bath. It may be necessary to
complete the drying in an oven at 105C.
For each filtrate, obtain the dry mass of salt and record it in your logbook.
Dry each sand sample in an oven until it is moisture-free (constant mass).
Record each sand mass in your logbook.
Results
Bulk
Supply
Mass of sand Mass of salt (g) Sand + salt % sand
(g)
total (g)
% salt
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Sampling details
Laboratory
samples
Mass
empty
beaker
of Mass
beaker
sample
of Mass of laboratory Mass of analytical
+ sample
sample used in test
1
1.
2.
3.
Analysis details for analytical samples
Mass
empty
paper
of Mass
filter paper
sand
of Mass
of Mass of basin Mass
check
on
+ empty
+ salt residue
cleaned out empty
evaporating
basin
basin
1
2
3
Analysis calculations
Mass of sand Mass of salt Recovered
Original mass of % sand
recovered (a) recovered
sand + salt analytical
Recovered
(b)
total (a) + (b) sample used in
test
% salt
Recovered
1
2
3
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Questions:
Discuss any differences between the composition of the original bulk supply and the test
results for your recovered samples.
Comment on your recovery check. (the agreement between the mass of each of your
analytical samples and the recovered sand + salt masses after the analysis. This tells how
reliable your results might be.
Suggest how you could improve the method to achieve better % composition and recovery
check agreement.
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Practical:
Date Completed:
Sampling Equipment
___________________
Teacher check _____________
Analyst signature ___________
Purpose
This practical is designed to familiarise the technician with equipment available to assist in
the taking of a valid sample.
Procedure
◗ Record the sample identity which you have been allocated by the tide zone from which
it was obtained.
◗ Using the riffles provided reduce the sample to approximately 100g and record the
analytical sample size in the table provided.
◗ Repeat step 2
◗ Select a nest of sieves and clean them thoroughly as demonstrated by the teacher, .
◗ Record the aperture sizes and assemble them so that the aperture decreases from
biggest at the top to the smallest next to the catch pan.
◗ Transfer one of the analytical samples to the top of the nest of sieves.
◗ Shake the sieves (with the lid on) for 5 minutes
◗ Using the A3 paper method demonstrated by the teacher, carefully capture and record
the mass of each fraction.
◗ Repeat the procedure with the other samples.
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Results:
Sample number
Sample mass
location on the beach from which the
sample was taken:
1
2
Sieve size
Mass sand
Sample 1
%
sand
fraction
in Mass sand
Sample 2
%
sand
fraction
in
Total mass
Questions:
Did you recover 100% of the initial sample? If not where did you gain or lose sample in the
method?
Did you have good agreement between your triplicate samples?
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Laboratory Techniques for Environmental Technicians
Chapter 8: Solution Preparation
General volumetric flask rules ( A recap)
◗ perform a sacrificial rinse using the intended solvent
◗ add the solute first, quantitatively, without loss
◗ add some solvent to ensure the solute dissolves completely
◗ mix well by inversion and swirling
◗ solvent is added until the bottom of the meniscus is on the line, mix well and support
the weight correctly
◗ label with name, concentration, date and appropriate safety advice.
1. Add solute
without loss
2. Dissolve in
some solvent
3. Invert and
swirl to mix
4. Make up to the
mark and repeat 3.
Steps to make up a solution in a volumetric flask
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Practical 5.4A Introductory Solution Preparation
Date Completed: ______________________
Teacher check ________________
Analyst signature ____________
Procedure:
You are required to prepare and validate known concentration solutions.
Using an analytical balance weigh out accurately the mass of each solute indicated in the
result sheet table.
Quantitatively transfer the solid to a 100mL volumetric flask and make up to the mark with
distilled water
The teacher will show you the checks to be made on your sample and the previously
prepared sample
Results:
Potassium Hydrogen Phthalate (Labelled as KHP) mass = 1.0 ±0.1g
Your Solution readings
Standard Sample readings
Sample mass:
Sample volume:
pH reading
Conductivity
Potassium chloride (labelled as KCl) mass = 5.0 ± 0.5g
Your Solution readings
Standard Sample readings
Sample mass:
Sample volume:
pH reading
Conductivity
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Questions:
Why is it important that an analytical balance is used to measure the mass?
Why is it important that the transfer of solid material is quantitative?
Why was a volumetric flasks specified in each case?
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Laboratory Techniques for Environmental Technicians
Chapter 9: Basic environmental laboratory testing
Laboratory Measurements
The following physical properties can be measured easily in the laboratory;
◗ pH
◗ refractive index
◗ density
◗ conductivity
◗ melting point
◗ boiling point
pH
Measurement of pH is a widely applied technique for monitoring agricultural, industrial and
biomedical processes. pH is a measure of the acidity of a sample, a low pH (ie <7) indicates
the sample is acidic, a high pH (>7) indicates the sample is basic.
pH may be measured using:
◗ indicator papers: paper impregnated with coloured dyes, whose colour is sensitive to
the pH of the solution to which they are exposed
◗ indicator solutions: solutions which are made up from the same coloured dyes which are
used for indicator papers. They are added directly to the test material where the
coloured form is used to assign a pH value.
pH meter: an electronic instrument which gives a direct readout of pH. It uses a glass probe
which generates an electrical signal in proportion to the hydrogen ion concentration of the
solution in which it is immersed. The pH meter must be calibrated prior to use to ensure the
validity of the result.
Practical work 7.1 pH measurements
Date Completed: ______________________
Teacher check ________________
Analyst signature ____________
Procedure:
The practical tasks provide experience with various detection methods for the
determination of pH. Your teacher will demonstrate the use of each method. Whilst working
with each consider the accuracy and efficiency of the method.
Indicator solutions, test papers and the pH meter are used to measure pH. Different
requirements for accuracy and speed will dictate which method is appropriate.
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Complete the table below by determining the pH of the samples provided, using the
methods as shown by your teacherResults:
Sample
Litmus
Red
Litmus BT
Blue
B
Initial
colour
Phen
olphtha
lein
Univer
sal
indicat
or
Othe
r?
Dip stick
pH
meter
Questions:
Which method do you consider the most reliable?
Which household chemical was the most acidic?
Which household chemical was the most alkaline?
What information does litmus paper give?
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Measurement of density
The density of any gas, liquid or solid is its mass per unit volume.
mass
Density=
volume
Mass is measured with a balance
Volume is measured by direct reading of a volume scale on the container or from
displacement measurements by immersion in a suitable liquid. The object is immersed in
the liquid and the apparent rise in level equals the volume of the object. Problems arise
when:
◗ the material is soluble
◗ the void spaces absorb fluid
◗ the material dissolves, swells, expands or shrinks
Liquid density measurements
Two basic methods are used
◗ density bottles (pycnometers): highly precise volumetric vessels which are filled with the
liquid in question and weighed.
◗ hydrometers: specially weighted floats which are immersed in the liquid.
Refractive index can be correlated with density. The producers and users of sugar-based
materials make extensive use of density as measured by hydrometry and refractometry to
analyse for sugar content.
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Practical:
7.7 Density
Date Completed: ______________________
Teacher check ________________
Analyst signature ____________
Procedure:
The practical is designed to give the student skills in determining the density of a variety of
different types of material, (irregular solids, granular solids, liquids etc) using a variety of
techniques (hydrometry, pycnometer etc).
Density of an irregular solid
◗ obtain three objects and record their identity
◗ measure and record the mass of the objects
◗ partially fill an appropriate measuring cylinder with water and record the initial volume
◗ Carefully immerse the object in the measuring cylinder
◗ record the new volume (final volume)
◗ Calculate the true volume
◗ complete the table by determining the density
Sample
Identity
or code
mass (g)
Initial
Volume
(mL)
Final
Volume
(mL)
True
Volume
(mL)
Density
(g/mL)
Literature
Density
You need to reference the source of your literature value data here.
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2. Bulk density of various granular solids
◗ obtain two different granular solids
◗ record the mass of an appropriately sized dry measuring cylinder
◗ add a predetermined volume of sample to the measuring cylinder and record this
volume in the table as the initial volume
◗ reweigh the measuring cylinder, record the mass
◗ gently tap the measuring cylinder to pack down the granular solid
◗ record the compacted (settled) volume
◗ add a known quantity of water to the measuring cylinder
◗ record the new volume in the measuring cylinder
◗ calculate the true volume of the granular solid
◗ complete the worksheet
Sample
I.D.
Empty
cylinder
mass (g)
Sample
+
Cylinder
Mass
sample
(g)
Initial
Sample
volume
Settled
volume
Volume
water
added
Final
volume
Mass (g)
Sample ID
True volume of Bulk Density
sample
Settled density
True Density
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3. Density of liquids
◗ weigh and record the mass of a clean, dry 25mL measuring cylinder
◗ add 25 mL of the liquid sample to the measuring cylinder
◗ reweigh the measuring cylinder and record the mass
◗ complete the worksheet
◗ follow the teacher’s instructions regarding disposal of the organic solvents
Sample
I.D.
Empty
cylinder
mass (g)
Sample + Sample
mass (g)
cylinder
mass (g)
Volume
sample
(mL)
Density
(g/mL)
Lit.
Density
Questions
How well do your practical results compare to the literature values?
What errors are present in the procedure you followed that would decrease your accuracy?
Identify the chemicals that would be of concern in the workplace.
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4. Density of liquids by hydrometry
◗ follow the teachers instructions regarding use of the hydrometer
◗ carefully float the appropriate hydrometer in a measuring cylinder of the allocated liquid
◗ read the value from the bottom of the meniscus
Sample
Hydrometer
Literature value
How well do your practical results compare to the literature values?
What errors are present in the procedure you followed that would decrease your accuracy?
Of the two methods (3 and 4) which do you prefer and why?
5.Determination of sugar content by hydrometry and refractometry
carefully measure the hydrometer reading of each of the sugar solutions provided, ensuring
that the hydrometer is dried before placement in a new solution
carefully measure the refractive index of each of the solutions following instructions given
by the teacher
determine the concentration of sugar in each of the unknowns by plotting
i) Hydrometer reading vs sugar content
ii) Refractive index vs sugar content
Sample identity
Hydrometer reading
Refractive index
5% sugar
10% sugar
15 % sugar
20% sugar
unk 1
unk2
You must plot a calibration graph (manually or by excel) for your RI and hydrometer results
and use it to calculate the sugar content of each of your unknowns.
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Laboratory Techniques for Environmental Technicians
Chapter 10: Gravimetric analysis
Gravimetric Analysis
Gravimetric analysis is the use of weighing to determine the amount of a component in your
sample.
Gravimetric analysis, or gravimetry is normally performed either as a :
◗ loss of volatiles procedure where the sample is heated to release volatiles such as
moisture or organic vapours and the change in mass is used to calculate the volatile
content, or
◗ precipitation or separation procedure where a component in the sample is isolated or
recovered in some form and this is weighed to complete the analysis.
Loss of volatiles gravimetric analysis
All loss of volatiles analysis is performed by a similar set of simple steps:
◗ weigh your fresh sample
◗ heat to remove volatiles
◗ reweigh to measure mass loss
◗ calculate percentage loss of mass or percentage residual mass
In practice, many problems can arise if you consider what might go wrong in each of the
steps. You need to:
◗ weigh enough sample to keep accuracy high but not too much so as to hinder vapour
loss
◗ heat but not for too long or too hot or too cold
◗ cool without additional losses or pick-up of volatiles
◗ reweigh without losses or gains
◗ be sure the container does not change weight during the procedure.
Heating to constant weight is the chief method used to ensure the container is not altered
during treatment. Empty containers are subjected to the same treatment as the sample
before any analysis and if their masses remain constant, then you can be sure they have
been heated to constant weight and their empty mass is correct.
Precipitation or separation gravimetric analysis
Precipitation or separation gravimetric analysis follows these simplified steps:
◗ weigh your fresh sample
◗ treat to isolate component of interest
◗ measure mass of pure, recovered material
◗ calculate percentage of component
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The treatment required to isolate the component can cover a wide range of techniques and
procedures including:
◗ digestion with powerful reagents such as enzymes, acids or oxidants
◗ buffering to ensure conditions are suitable for recovery of the analyte
◗ addition of chemicals designed to selectively and quantitatively capture the component
of interest.
◗ use of prolonged heating or contact with solvents such as occurs with Soxhlet extraction
or Dean and Stark moisture determination.
Validation of Gravimetric Analysis
A requirement with all laboratory testing is some evidence or assurance of the reliability of
your answer. This evidence comes in four major ways:
◗ replicate analysis
◗ control or standard analysis
◗ literature values
◗ alternative testing procedure
Common containers for gravimetry
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Practical: Analysis of a mercury chloride and determination of its
empirical formula.
Date Completed:
___________________
Teacher check _____________
Analyst signature __________
Introduction
In this practical you will be given a compound of mercury and chlorine and by decomposing
the compound into its elements you will be able to calculate the percentage composition
and the empirical formula of the original compound.
Decomposition of the mercury chloride is achieved by heating the substance with a mixture
of sodium hypophosphite and hydrochloric acid until the element mercury is produced.
WARNING
All compounds of mercury should be considered to be highly poisonous- don’t breathe in
the fumes from the solution and make sure that the mercury is covered by solution during
heating. Wash your hands thoroughly at the completion of the exercise and make sure all
spillages are cleaned up correctly!
Procedure:
◗ Obtain a sample tube containing a mercury chloride and record the number on the tube.
◗ Weigh a clean, dry evaporating basin, labelled with your name. Also weigh the sample
tube plus contents
◗ Carefully add the mercury chloride to the evaporating basin.
◗ Collect 20 mL of distilled water and use some to rinse the contents of the sample tube
into the evaporating basin. Add the remainder of the water to the evaporating basin. Reweigh the empty, dry sample tube.
◗ Add 10 mL of sodium hypophosphite/hydrochloric acid solution from the dispenser
bottle to the contents of the evaporating basin.
◗ Place over a steam bath in a fume cupboard and stir to mix the reactants (Care - liquid is
corrosive)
◗ Continue heating and stirring until mercury collects in large silvery globules. (Avoid
inhaling vapour while heating)
◗ When all the mercury globules have coalesced, remove the basin from the steam bath
and decant the colourless liquid AS DEMONSTRATED BY THE TEACHER.
◗ Wash the contents of the basin twice with water and once with small amounts of
methylated spirits. Collect all washings in the waste container provided.
◗ Dry the mercury by pulling pieces of filter paper through it as demonstrated.
◗ Re-weigh the basin containing the dry mercury.
◗ Return the mercury to the class collection bottle.
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Results
Sample ID
Mass evaporating basin
Mass tube + contents
Mass empty tube (after washing)
Mass evaporating basin + mercury
Calculations
Calculate by subtraction of appropriate measurements ◗ Mass of mercury chloride used
◗ Mass of mercury metal produced
◗ Mass of chlorine originally present
From these figures calculate
a) the percentage composition of your compound
b) the empirical formula of your compound
c) suggest a suitable name for your compound
Discussion
Comment on any errors in your analysis and how this may be overcome.
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Practical: Gravimetric determination of sulfate in bore water 8.6
Date Completed:
___________________
Teacher check _____________
Analyst signature __________
Purpose
To determine the sulphate content of bore water using a precipitation method.
Procedure
◗ Accurately weigh approximately 2.og of fertiliser and transfer to a 500mL beaker. Record
the mass labelled on the beaker in the Results section
◗ Add approx 50 mL 0.1M hydrochloric acid from the dispenser
◗ Gently heat the beaker to boiling on a hot plate to dissolve the fertilizer (there will be
some undissolved material).
◗ While the heating is proceeding –
◗
From the desiccator weigh a #4 sintered glass crucible. Do NOT touch the glass surface
with bare hands, use tongs to avoid contamination. Record the Empty mass in the result
table.
◗ Assemble a vacuum filtration set up utilizing a Buchner funnel (fit with drip tip towards
the vacuum spout) and an appropriate sized type 54 filter paper.
◗ When the fertilizer solution boils remove the beaker from the heat and allow it to cool
before filtering. (Use an ice bath to cool if necessary).
◗ Filter the fertilizer solution. Rinse the beaker 3 times with distilled water to ensure all
the solution passes into the vacuum flask.
◗ Discard the filtered-out insolubles (on the filter paper) and return the filtrate (the liquid)
to the original labelled beaker, again using 3 small distilled water rinses.
◗ Add 25 mL (by measuring cylinder) of Barium Chloride to the beaker. The reaction
converts the soluble sulfate ions to insoluble Barium Sulfate.
Ba2+(aq) + SO42+(aq) 
BaSO4(s)
◗ Return the beaker to the hot plate, cover with a large watch glass and heat gently (no
boiling) for about 1 hour, then allow to cool.
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Fit the pre-weighed sintered glass crucible to your vacuum filtration set up. (No touching
with fingers!)
Use the sintered glass crucible to filter the
Barium Sulfate crystals.
After transferring all the contents of the beaker
wash the crystals 3 times with about 10 mL of
purified water each time.
Wash the inner wall of the sintered glass
crucible and the crystals with ethanol.
Leave under vacuum for about 1 minute, this
will ‘vacuum dry’ the crucible and crystals.
Transfer the crucible into a 110° drying oven.
Weigh the crucible with its load of dry BaSO4
crystals (no touching!). Record the mass in the
table .
Results
Sample 1
Sample 2
Crucible ID
Mass of empty crucible (g)
Mass of empty crucible (g) after
heating to constant weight
Mass of crucible + residue (g)
Mass of crucible + residue (g) at
constant weight
Mass of residue (g)
Average mass of residue (g)
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Calculations
The value reported is typically in the units of mg sulphate / L
Sulfate = average mass of residue x 40 x 0.4113 x 1000
Questions
What is a supernatant solution?
What is the purpose of adding more BaCl2 to the supernatant fertilizer solution?
What is happening to the crystals while they are kept in the hot solution for an hour?
Why were the Barium Sulfate crystals washed repeatedly?
Barium Chloride crystals have 2 molecules of water associated with each molecule of BaCl2
ie. BaCl2•2H2O How does this affect the way you would prepare a 5% solution of BaCl2?
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Practical: Experimental Investigation of combustion of magnesium
An open-ended investigation is one in which the student is asked to investigate a particular
problem, with no specific information being provided about the end result. The student is
asked to consider the possibilities and make judgements on how to proceed. In this practical
you are to investigate the combustion of magnesium. A method will be provided but before
commencement you are being asked to decide what will happen to the magnesium, will it
change in any way, will it look the same, how might it change, will it be the same after as it
was before……
Write your answers to the above before commencement of the practical
Combustion of magnesium
When magnesium is burnt (combusted) in air, it reacts with the oxygen gas forming a white
ash of magnesium oxide. This is a highly exothermic reaction producing a large amount of
energy as heat and light.
Equation
Magnesium(s) + oxygen(s)  Magnesium oxide(s) + energy
This practical investigates the mass of magnesium compared to the mass of the magnesium
oxide (ash) produced. In an enclosed vessel (crucible) the magnesium will quickly use up
available oxygen as it burns and so we will need to open the lid several times to allow air to
enter. If sufficient oxygen is not present the nitrogen (from the air) will often combine with
the magnesium forming white magnesium nitride and so the ash contains both magnesium
compounds.
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Date Completed:
___________________ Teacher check _____________
Purpose
Analyst signature __________
To determine how the mass of magnesium burnt compares with the mass of the ash formed
Procedure
Weigh an empty clean crucible with its lid and record the mass
Place three 2 cm pieces of magnesium ribbon in the empty crucible and record the mass
Heat the crucible, with the lid slightly ajar, in a strong Bunsen flame
After several minutes check that the magnesium has been burnt.
Allow the crucible to cool to room temperature with the lid on.
Weigh the crucible with contents and lid and record the mass
Results
Mass of crucible + lid
=
Mass of crucible + lid + magnesium
=
Mass of magnesium
Mass of crucible + lid + ash
=
=
Mass of ash
=
Conclusions
The law of conservation of mass (matter) states that in any chemical reaction the total mass
of reactants is equal to the total mass of products since the reaction involves just the rearranging of the atoms which are already present
“Matter can neither be created nor destroyed, it can only be changed form one form into
another”
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Questions
Assuming that your magnesium totally reacted to produce magnesium oxide, what mass of
oxygen must have been used from the air?
When a piece of wood burns the resulting ash produced weighs much less than the original
wood. Where did all the other atoms go?
The main elements in wood are carbon, hydrogen and oxygen. What compounds do you
think carbon and hydrogen would form when the wood is burnt in air?
Carbon would combine with oxygen to form ……………………………….
Hydrogen would combine with oxygen to form …………………………….
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Practical: Gravimetric determination of Ni by precipitation with
dimethylglyoxime.
This method involves converting nickel ions which are soluble in water into a complex which
has reduced solubility at an increased pH. The procedure involves the use of a complexing
agent dimethylglyoxime and strict pH monitoring to form an insoluble solid. The complex
formed is Nickel dimethylglyoxime. The analysis is gravimetric as the precipitate formed is
collected and weighed to determine the mass of nickel.
Date Completed:
___________________ Teacher check _____________
Purpose
Analyst signature __________
To determine the nickel content of a water solution
SAFETY
The technique involves digesting the solution over a steam bath. Remember steam and skin
do not go well together.
Ammonia has a strong odour and work with ammonia should be carried out in the fume
hood. When smelling do not inhale deeply (Use the wafting technique as demonstrated by
your teacher.)
Concentrated HCl and skin to not go well together.
The sintered glass crucibles will be hot when removed from the drying oven
Procedure
Clean two sintered glass filter crucibles with dilute HCl and dry in a 110 oC oven.
Pipette out two 20 mL samples of the supplied unknown into two 600 mL beakers
Dilute to about 150 mL with purified water
Add approximately 3g of citric acid to each solution and stir to dissolve. Leave the glass
stirring rod in the beaker
Slide a piece of red litmus paper down the side of the beaker until it is mostly immersed in
the solution, but sticks to the side.
Add 5 M ammonia slowly until the litmus turns blue, and then add concentrated HCl slowly
until the paper returns to red. Add 2 mL more of concentrated HCl. Remove the litmus
paper.
Heat the beakers on a hotplate to about 80oC but do not allow to boil.
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Remove the beakers from the hot plate and slowly add 50 mL 1% dimethylglyoxime (DMG)
solution with stirring. If a red precipitate begins to form, add concentrated HCl dropwise
until it redissolves, and then continue to add the DMG solution.
Add 5M ammonia slowly with stirring until a red precipitate forms and the solution smells
strongly of ammonia.
Stir well and stand on a steam bath for 30 minutes.
Check that the solution still smells of ammonia, (add extra if necessary)
Add 2 mL extra DMG using a plastic pipette, and stir well.
Cool the beakers to room temperature and filter the solutions through the previously dried
and tared sintered glass filter funnels.
Wash the precipitate in the crucible 3 to times with water
Finally wash the precipitate once with 20-30 mL of 30% aqueous ethanol, which dissolves
any remaining DMG from the precipitate.
Place the crucibles in an oven at 110oC for at least 2 hours.
The crucibles should be cooled in a desiccator and then reweighed.
Results
Sample 1
Sample 2
Volume solution taken
Mass of
crucible
sintered
glass
Mass crucible + Ni(DMG)2
Calculations
1. Ni2+(aq) + 2DMG

Ni(DMG)2(s)
Moles Ni2+ in original solution
=
mass Ni(DMG)2
= moles of Ni(DMG)2 formed
.
288.88
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Mass nickel in original solution = moles of nickel x Formula mass nickel
Sample 1
=
Sample 2
=
Ave
=
%w/v of nickel in the sample = mass of nickel in sample x 100
vol of sample
QUESTIONS
Why was it important to limit the amount of nickel in the sample by only taking 20 mL?
Why was the precipitate washed in aqueous ethanol?
Why was a sintered glass crucible used rather than filtration through paper?
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Chapter 11: Volumetric Analysis
This section builds on the chemistry in chapter 2 and the theory is provided only as a means
of illustrating the concept. It does not form an examinable part of the units being
undertaken in the intensive session. It is the titration skill that will need to meet the
required level of competence.
The titration process
The terms volumetric analysis, titrimetry and titration are used interchangeably to describe
a procedure which analyses chemicals in solution by accurate volume measurement.
Principles of titration
Titration is used to determine the concentration of an analyte in solution (the unknown
solution) and hence the components of any material which can be made up into solution
may be analysed this way. The method operates by reacting the analyte in the unknown
solution with another solution (the standard solution). The process is carried out so that
reaction between the two can be stopped exactly at the point when the last trace of analyte
has been used up. The volume of both standard and unknown are used to obtain the
required quantitative information.
To perform titrimetry, the following need to be available:
◗ standard solutions
◗ volumetric glassware
◗ a suitable method of detecting the end-point
◗ a minimum quantity of analyte
◗ a standard that reacts appropriately with your analyte
The burette is filled with a known
concentration of chemical R which
destroys A
An accurately measured amount
of
the material containing analyte
(A)
R is added
until the very
last of A is
destroyed –
but not one
drop more of
R is allowed
An aliquot of the
is put into a known volume of
test solution is
solution
now measured
for titrimetric
analysis
Only A
This flask now contains a fixed
number of particles of A
A is being
consumed
The amount of A
is calculated
from the
amount of R
which
was added.
Both
A and R
no longer exist
by R as it is
added
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The practical term, end point, is used to describe the volume of standard solution which was
used to just consume the aliquot of analyte. The theoretical term, equivalence point, is used
to describe the volume needed to provide the exact molar quantity for the analyte to react
completely with the standard. The end point is usually a little more because you need to
have a little extra to cause the indicator to change colour.
Theoretical Requirements
Titrations require the following conditions:
a rapid and complete chemical reaction between the analyte and the standard. Those
reactions which are sufficiently fast and complete to be suitable for volumetric analysis
include acid-base, redox, complexation and precipitation reactions. This enables an
extraordinary number of analytes to be tested by titration and this versatility makes the
procedure so universally popular. Some titrations are carried out at higher temperatures to
ensure the reaction is rapid and complete.
a balanced equation for the reaction. To be able to calculate the final answer, you must
know the molar ratio of reaction between the analyte and the standard and that the
chemical change has gone to completion.
a means for detecting the disappearance of the last trace of analyte. Because the volume of
addition is judged by the disappearance of analyte, a signal which is linked to this event is
mandatory. This signal can be a colour change which is detected by eye (or some other
property such as a temperature or conductivity change, which is detected instrumentally).
The most common methods use indicator solutions which are coloured dyes with two
coloured forms — one colour exists while analyte is still present and the other colour is
created when the analyte disappears. The analyst performing the procedure needs to be
alert to detect the first permanent colour change. The addition is stopped and the volume of
titrant (solution added from a burette) is measured. The titration is said to be at its end
point and this volume becomes one of your measured end-point volumes.
Practical Requirements
To perform titrimetry, the following need to be available:
standard solutions (solutions of exactly known concentration). These should have a
reasonable shelf life and be easy to prepare from ingredients that are convenient to handle.
A more stringent set of requirements appears in Section 9.3 on types of standards.
volumetric glassware (burette, pipettes and volumetric flasks) and support equipment
a suitable method of detecting the end point
a minimum quantity of analyte. Below this minimum, the method will not be able to detect
the presence and then the absence of analyte reliably. Typically, titration can not reliably
measure concentrations of less than 0.1%, unless some form of pre-concentration is used
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a standard that reacts appropriately with your analyte. You can’t just take any standard
solution and do a titration and expect a valid result. There must be a known reaction and
indicator that detects a true end point. Not all analytes have appropriate standards for
titrimetry. Two analytes for which titrimetry does not work are nitrate ions, and fat in food.
Selection of Indicators
Indicators are coloured chemicals which have two forms — each with a different colour. The
chemical form and colour of the dye are affected by the chemical conditions in the solution
in which it is dissolved. Ideally, an indicator is selected because the presence of the analyte
favours one coloured form and the absence of analyte favours the other. From the practical
viewpoint, catalogues of indicators reveal there are scores of choices and so a selection
procedure is required. Initially you might suspect that each analyte requires its own
indicator and any indicator can only identify one analyte, but this is not true. The choice of
the correct indicator is determined by the conditions at the end point or equivalence point.
Indicators are classified as:
◗ acid-base
◗ redox
◗ adsorbtion
◗ or complexometric and must match the type of reaction occurring during titration.
All indicators have precisely defined conditions for changing colours. For example, the key
factor for acid-base indicators is pH at the end point. Table 9.1 shows some common acidbase indicators, their colours, and their pH changeover range. The names of indicators are
deceptively simple and easily confused (e.g. methyl violet is not even violet). Many have
similar colours (e.g. yellow and red), but change at different pH intervals. The direction of
change may also be reversed (e.g. methyl orange is red at low pH and yellow otherwise,
whereas phenol red is yellow at low pH and red otherwise).
Acid–base indicators
Indicator
pH
changeover high pH colour
range
low pH colour
Methyl violet
0.0–1.6
blue
yellow
Methyl orange
3.2–4.4
yellow
red
Phenol red
6.6–8.0
red
yellow
Bromothymol blue
6.0–7.6
blue
yellow
Phenolphthalein
8.2–10
pink
colourless
Alizarin yellow R
10–12
red
yellow
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End point pH is determined by the strength (not concentration) of acid and base used, and
the natural pH at the equivalence point when exact molar proportions are present for all the
acid to neutralise all the base.
Weak acid–strong base equivalence points (e.g. ethanoic acid–sodium hydroxide) are all
greater than pH 7.
Strong acid–weak base equivalence points (e.g. HCl–Na2CO3) are all less than pH 7.
Your procedure needs to have information on end-point pH or a specified indicator.
The Titration Procedure
The diagram below shows the general equipment used for titration. A burette is used to
contain one of the solutions (either the standard or the analyte solution) and a conical flask
or conical beaker holds the other. The following steps are then performed.
Burette
Rinse the burette with a small amount of the fill solution as the final step before filling.
Obviously if there is any doubt, a more thorough cleaning will be required prior to the final
filling step. (Remember the ‘magic number’ three)
Overfill with the required solution.
Check there are no air bubbles above or below the tap and in the bore of the tap.
Wipe any suspended liquid off the burette tip
Zero the burette or record its initial volume. Re-wipe the burette tip.
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The indicator
is added and
you are ready
to titrate
An aliquot of
the other
solution is
placed in the
titration flask.
(e.g. the
analyte)
Titration
flask
One solution is used to fill the burette (e.g.
the standard)
Apparatus used to perform titration
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Titration Flask
Rinse the flask with distilled water as the final step before adding the aliquot of solution. It
may be left wet.
Rinse a pipette with a small amount of the fill solution as the final step before filling.
Obviously if there is any doubt about contamination, a more thorough cleaning will be
required.
Add the aliquot of solution. Sometimes an accurately weighed primary standard solid is used
as the reactant, which needs a suitable amount of water added to dissolve it, prior to
additions from the burette.
Add a suitable indicator.
Rough titration
Add the solution from the burette in a slow steady stream to the titration flask while
swirling vigorously.
Observe the colour of the indicator carefully for some change in its shade or hue. This is a
sign of an approaching end point.
Stop the addition when you see the sudden colour change. Record this volume reading
which is a rough end point. Obviously, with the first titration you will go many drops even
many mL past the true end point but the purpose was to get a rough idea quickly rather
than spend a lot of time adding single drops, mixing and waiting to see what happens. A 25
mL titration is about 500 drops and would be very slow and painful to do dropwise!
Accurate titration
Add the bulk of rough titration volume (say 80–95%) rapidly.
Rinse the walls of the flask with distilled water to ensure all reactive species are in the
mixing part of the flask This is essential when you are only few drops short of the end point.
Complete the addition dropwise so that the end point is overshot by no more than one or
two drops. Any more than this means you will need to reject this measurement.
Repeat the titration procedure with fresh materials until you have an acceptable set of end
point volumes (at least in triplicate with less than 0.2 mL range between all results).
Calculate your mean, range, and absolute and relative precision. Titration results are only
acceptable if they lie within a specified precision range (at least within  2% but  0.2% is
possible with some standard methods).
Some indicators fade on standing and so it will appear that your end point has changed but
the general rule is to accept the first change which is permanent for 10–20 seconds.
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Ideal titration
Ideally, the addition of a fraction of a drop (< 0.05 mL — the limit of reading of a burette),
will produce a permanent colour change different from the previous colour. The volume
added at this point is the true end point. (The equivalence point, which describes the
volume added to provide the exact molar amounts of analyte and standard at the end point,
can be used to calculate the ideal end point volume but it does not always match exactly
what is observed in practice and you don’t usually know the answer anyway).
Replicate determinations in an ideal analysis will all lie within 0.05 mL. The accumulated
errors of the glassware used will produce a range of readings of at least this magnitude. All
novices should aspire to performance which approaches these standards.
Types of Standards used in Titration
Standard solutions must have the following attributes:
◗ an accurately known concentration
◗ contain an active chemical species which reacts stoichiometrically with analytes of
interest (i.e. you know the ratio of moles at the endpoint and the fact that the reaction
is complete)
◗ reasonable shelf life in terms of maintaining a known concentration
◗ belong to a known chemical reaction class (e.g. acid or base, oxidant or reductant,
chelating agent or complexible metal ion), so that correct matching of analyte and
standard occurs.
Two types of standard solutions are available in laboratories:
primary standard solutions made up from primary standards by accurate weighing and
dilution to volume in quantitative glassware. A primary standard is normally a solid chemical
of precisely known purity (usually 99.9% or better) which is easy to weigh, tolerates
exposure to the atmosphere and has a reasonable shelf life after opening the container.
Naturally, it must have a reactant concentration which can be calculated from the original
mass. Only very few chemicals qualify as primary standards, but examples exist for each of
the major classes of reactions for which titration is used. Commonly, sodium carbonate solid
is a primary standard alkali and potassium hydrogen phthalate solid is a primary standard
acid.
secondary standard solutions made up from secondary standards by approximate weighing
and dilution to volume to give an approximate concentration. The secondary standard is
defective in some way which prevents it being suitable for precise weighing. Generally it is
because of unreliable purity, problems with exposure to the air or limited shelf life. Thus a
solution containing the chemical species of interest can be made up, but with only an
approximately known concentration. Hydrochloric, sulfuric, nitric and ethanoic (acetic) acids
are all only suitable for use as secondary standard acids in their normal forms. Sodium
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hydroxide, potassium hydroxide, barium hydroxide, and ammonia solution are only suitable
for use as secondary standard alkalis in their normal forms. They must all then be
standardised (e.g. titration with a primary standard) to establish the true concentration.
Calculations
Many methods exist to process titration data into analyte concentration and other
information on composition. One method which restricts itself to moles and molarity is
described here.
All quantities which describe the analyte and the standard need to be converted to moles.
All concentrations must be converted to molarity measured in moles per litre.
Commonly used terms and their symbols and their units are given in the following table
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Term
Symbol
Unit
moles
n
mol
molarity
M
mol. L–1 (mol/L)
volume
V
mL but for some calculations it must be
L
mass
m
g
formula weight
FW
g. mol–1
The formulae used to convert titration data to analyte data are found in below.
Combination
Formula
moles of solute in n = M x V
solution
Equation
No.
Conditions
1
V must be in L
moles of a solid
n = m  FW
2
molarity from moles
M= n  V
3
V must be in L
molarity from mass
M = (m  FW) 4
 V
V must be in L
Examples
Moles and molarity calculations
(a)
How many moles of H+ in 25.0 mL of 0.1 M HCl ?
data: M = 0.1 mol. L–1 and
V = 25 mL which must be converted to L = 25 x 10–3 L
Use Equation 1 to give n = 0.1 x 25 x 10–3 moles of HCl
ANS.
(b)
n = 2.5 x 10-3 moles of Hcl and also of H+
How many moles of CO32– in 0.40 g of Na2CO3?
data: FW = 23 + 23 + 12 + 16 + 16 + 16 = 106 g/mol
m = 0.40 g
Use Equation 2 to give n = 0.4  106 moles of Na2CO3
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ANS.
(c)
n = 3.77 x 10–3 moles of Na2CO3 and also of CO32–
What is the molarity of 25 mL of solution containing 2.5 x 10 –3 moles of Hcl ?
data: V = 25 mL which must be converted to L = 25 x 10–3 L
n = 2.5 x 10–3 moles
Use Equation 3 to give 2.5 x 10–3  25 x 10–3
ANS.
(d)
M = 0.1 mol. L–1 of Hcl (or 0.1 M HCl)
What is the molarity of 100 mL of solution containing 0.40 g of Na2CO3?
data: V = 100 mL which must be converted to L = 100 x 10–3 L
m = 0.40 g
FW = 23 + 23 + 12 + 16 + 16 + 16 = 106 g/mol
Use Equation 4 to give n = (0.4  106)  (100 x 10–3)
ANS.
n = 0.0377 mol. L–1 of Na2CO3 (or 0.0377 M Na2CO3)
Mole Ratios in Titration Reactions
Not all titration reactions proceed in simple 1:1 ratios. The reaction ratio must be supplied,
as in the case of standard methods, or checked by writing the balanced equation.
Titration Calculations
Calculate the number of moles in the volume of standard solution or in the mass of primary
standard solid used.
Determine the
supplied values.
reaction
ratio
from
the
balanced
equation
Calculate number of moles of the analyte which would
number of moles of standard from step 1. Use the ratio from step 2.
or
use
the
react
with
the
Use the volume of analyte solution to calculate its concentration.
Examples
Titration calculations
1.
25.0 mL of an unknown H2SO4 solution is titrated to an end point of 21.3 mL with
0.1034 M NaOH. Calculate its concentration.
Step 1 Number of moles of NaOH = 0.1034 x 21.3  1000 = 2.20 x 10–3 mol.
Step 2 H2SO4 + 2 NaOH
Na2SO4 + 2H2O.
Therefore, 1 mole of sulfuric acid reacts with 2 moles of NaOH,
or 1 mole of NaOH is equivalent to 0.5 mole of sulfuric acid.
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or for every mole of NaOH consumed there was 0.5 mole of H2SO4 present.
Step 3 Number of moles of H2SO4 = 0.5 x moles of NaOH = 1.10 x 10–3 mol.
Step 4 M = 1.10 x 10–3  25 x 10–3 = 0.0440 mol. L–1.
The concentration of H2SO4 was found to be 0.0440 mol. L–1 (0.0440 M
H2SO4).
2.
19.4 mL of HCl solution is needed to obtain an end point with 0.153 g of Na 2CO3.
Calculate the concentration of the HCl.
Step 1 Number of moles of Na2CO3 = 0.153 106 = 1.44 x 10–3 mol.
Step 2 2 HCl + Na2CO3
2 NaCl + H2O + CO2.
Therefore, 2 moles of HCl react with 1 mole of Na2CO3,
or 1 mole of HCl is equivalent to 0.5 moles of Na2CO3.
or for every mole of Na2CO3 consumed there was 2 moles of HCl present.
Step 3
Number of moles of HCl = 2 x moles of Na2CO3 = 2.88 x 10–3 mol.
Step 4 M = 2.88 x 10–3  19.4 x 10–3 = 0.149 mol. L–1.
The concentration of HCl was found to be 0.149 mol. L–1 (0.149 M HCl).
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Practical: Practice titration
Date Completed:
___________________ Teacher check _____________
Purpose
Analyst signature __________
Results:
1.
Observation of endpoint colour
Colour of solution Colour of solution Colour of solution
using
screened using
Using M.O.
M.O.
phenolphthalein
Acid HCI
Base NaOH
End point
2.
Practice Titration
Concentration of supplied acid
Volume of alkali pipetted into titration
flask
Volume of acid needed for titration
end point
1
2
3
4
Average volume acid used for end
point
3.
Calculation of concentration of supplied alkali
Concentration alkali =
[Concentration acid
x Volume acid] / Volume base
=
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Questions:
What are the possible sources of error in the use of a burette?
Why should all readings be written down and checked before draining or refilling burettes?
Place each of the following pieces of glassware which may be used in a titration in the
appropriate preparation column.
Pipettes, burettes, volumetric flasks, titration conical flasks, beakers and bottles used to
hold standard and other solutions.
Used dry
rinsed with solution to rinsed
be placed in it
water.
with
distilled
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Practical: Ethanoic acid content of Vinegar
Date Completed:
___________________ Teacher check _____________
Purpose
Analyst signature __________
To determine the ethanoic acid content of vinegar.
Chemistry:
Complete the equation to enable the mole ratio to be determined
Sodium hydroxide + ethanoic acid  sodium ethanoate + water
NaOH(aq) + CH3COOH(aq)
 NaCH3COO(aq) + H2O(l)
Results:
Volume of original vinegar taken (mL)
=
Volume of diluted vinegar prepared (mL) =
Vol. of diluted vinegar titrated (mL)
=
Concentration of NaOH used
=
Titration volumes
Average titration value
=
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Calculations:
moles of NaOH used
= Concentration NaOH x Average titration / 1000
=
moles ethanoic acid
= moles of NaOH
(because the mole ratio = 1:1)
=
Conc diluted ethanoic acid
= moles ethanoic acid / (volume of aliquot/1000)
=
Conc original ethanoic acid
= Concn of diluted vinegar x 10
(original vinegar solution)
=
Mass (g) ethanoic acid per L
= Concn original ethanoic acid x 60
=
% w/v ethanoic acid in vinegar
= mass ethanoic acid per L / 10
=
Questions:
How would you titrate if brown vinegar had been used in place of white vinegar?
How does your answer for the vinegar compared to food authority standards for legal
vinegar? (The ethanoic acid content must be at least 4.0% w/v).
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Practical: Sodium Carbonate Content of Washing Soda
Date Completed:
___________________ Teacher check _____________
Purpose
Analyst signature __________
To determine the sodium carbonate content of washing soda by titration
Equation:
Sodium carbonate + hydrochloric acid 
Results:
Mass of washing soda: (approx 3.6 g)
=
Size of volumetric flask:
=
Aliquot volume (mL)
=
Aliquot volume (L)
=
Concentration of HCl
=
Titration volumes (mL)
Average titration (mL)
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Calculations:
moles of HCl
= concn of HCl x average titration volume / 1000
=
moles Na2CO3
= moles of HCl / 2
in aliquot
=
Concn Na2CO3
= moles Na2CO3(dil)  aliquot volume(L)
=
Mass Na2CO3 in 250 mL = (Concn Na2CO3  4) x 106
volumetric flask
=
% sodium carbonate
= mass sodium carbonate in 250 mL x100
Mass of initial washing soda
=
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Practical: Chloride by titration with silver nitrate
Date Completed:
___________________ Teacher check _____________
Purpose
Analyst signature __________
To determine the salt content of potato chips
Sodium is an essential element in the body – playing a major role in the functioning of
nerves and the electrolyte balance in the blood. It also is an important substance in many
other organisms and is partly responsible for the taste (or lack of it) in natural waters. When
combined with chloride ions, sodium imparts a salty taste to water. Too high a
concentration of these ions leads to the condition known as salinity, which destroys soil
making it unsuitable for crop growing and general agricultural use. In drinking waters small
amounts of sodium and chloride are actually desirable to help give the water taste. Water
should be monitored to ensure that the levels remain acceptable for consumers.
Chloride ions are relatively easy to monitor, as they may be analysed using a simple titration
with silver nitrate.
Procedure
Weigh accurately about 3 g of chips (in triplicate) into 250 mL conical flasks
Add about 50 mL of purified water and a small amount of calcium carbonate CaCO 3 powder
and swirl for about 3 minutes
Add about 1 mL of chromate indicator
Titrate with standardised 0.1M silver nitrate until endpoint.
Record your results and calculate the amount of salt in the potato chips
Results
1
2
3
Mass of potato chips (g)
Conc AgNO3
Volume AgNO3(mL)
Calculations
Balance the equation
AgNO3 + NaCl 
AgCl(s) + NaNO3
Volume (L) Titrant
=
volume AgNO3 / 1000
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=
Moles Ag+
=
concentration AgNO3 x volume AgNO3 (L)
=
moles Ag+
=
Moles chloride ion
=
Moles sodium chloride = moles chloride ion
=
Mass sodium chloride = moles chloride x FWt(NaCl)
=
% NaCl
= mass sodium chloride / mass chips x 100
=
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Practical: Determination of hardness in water
Date Completed:
___________________ Teacher check _____________
Purpose
Analyst signature __________
To determine the hardness of a supplied water sample.
Water Hardness
Water hardness is a common problem in outback Australia and areas whose water supply is
derived from a river source. Hardness is defined as “the ability of water to cause
precipitation of insoluble salts of higher fatty acids from soap solutions”. In other words
water is hard when it won’t lather with soap! It is generally caused by high concentrations
of Ca2+ and Mg2+ being present in the water. When used in water systems both of these
hard ions (Ca2+ and Mg2+) play a role in forming coatings on the insides of pipes.
In this practical you will determine the total hardness of a water sample using the EDTA
titrimetric method. When reporting the result it is expected that the method used for
determination will be quoted e.g. hardness (EDTA).
Water hardness is removed chemically by the addition of sodium carbonate which causes
the precipitation of Ca2+ and Mg2+ from solution. Another procedure (industrial use) utilises
ion exchange resins called zeolites, which adsorb the divalent cations as they are pumped
through them. This is important for reducing the level of boiler scale (precipitation caused
by hard water) in pipes in industrial processing plants.
Typical concentrations of ions (expressed as CaCO3 equivalents) – [from G. Laidler, Environmental Chemistry]
Concentration
Description of
CaCO3 mg/L
Water Quality
0-50
soft
50-100
moderately soft
100-200
slightly hard
200-300
hard
> 300
very hard
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Procedure
Pipette a triplicate aliquot of the supplied water sample into 3 separate conical flasks. Add
enough purified water to bring the volumes to about 50 mL.
Add 5 mL of ammonium/ammonium chloride buffer and about 50 mg of Eriochrome black T
indicator (your teacher will demonstrate the correct amount of powder) to each flask.
Use the burette containing 0.01M EDTA solution, and titrate to the end point, which is the
first permanent appearance of a blue colour. Note: the colour change is reasonably slow, so
your first titration will probably overshoot.
Repeat steps 2 - 3 until to you have three titration values within a 0.5 mL range. Record your
best 3 titrations on the result sheet
Complete the calculations to determine the hardness of your water sample. After your
calculation compare your value with a table provided and report the relative hardness of
your sample.
Results
Molarity of EDTA
Titration Values
M
Titration 1 =
mL
Titration 2 =
mL
Titration 3 =
mL
Ave
mL
=
Calculations
A.
Moles of CaCO3
=
molarity EDTA x titration (mL)
1000
=
B.
Mass of CaCO3 (FWt 100)
=
A x 100
=
Water Hardness (as CaCO3)
=
B x 1000000
mL of sample
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=
Water Hardness mg/L
=
mg/L
Hardness Rating
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Chapter 12: The Open Ended Problem
Open ended project
You have been working in your chemistry class and laboratory class towards an
understanding of basic concepts in both the laboratory and chemistry,
This task is designed to enable you to show your competence in both areas.
You are to devise a simple laboratory project that will allow you to separate a mixture of
sand, salt and sawdust and provide the percentage composition of the original mixture. You
will be guided in class by your laboratory teacher and also you chemistry teacher.
Requirements:
Submission to your teacher of a flow chart showing the general laboratory procedure
Identification of the physical And/or chemical properties you will use to make the
separation
A completed laboratory request form for each practical session
Completed 5 minute risk assessments for each practical session
A work journal indicating how your ideas changed each session
A log book that records your data in a scientific manner
A written report following the format provided by your teacher
You will be required to submit a report at the conclusion of the project.
The following questions could be useful as you design your project:
What properties of each component could be useful for separation?
What equipment will you require for the separation? Is it readily available in the laboratory?
What hazards exist in your separation procedures?
What SDS will you require?
Do you need to know the original mass of the mixture?
How many sessions do you need for the separation?
Are there other methods that could work? Why did you choose a particular method?
Why is it important to have a logbook for your data?
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Chapter 13: Appendix/Assessment information and completion record
Assessment 1
For the basic Chemistry component there are 4 assignments to be completed
Assessment 2
For the calculation section there are 2 assignments to be completed
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Assessment 1
Assignment 1
a) Explain how you could separate the following mixtures in order to get the maximum yield
of the highlighted component
vegetable oil and water
talcum powder and sugar
iron filings and sand
b) Give the symbol for the following elements
Gold
Helium
Tungsten
Sodium
Nitrogen
Calcium
Iron
Chlorine
Molybdenum
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Assignment 2
Complete the following chart
Name
Formula
Name
Formula
potassium fluoride
dinitrogen trioxide
carbon disulfide
phosphorus pentachloride
calcium phosphate
sodium
carbonate
silver nitrate
lead (II) acetate
ammonium sulfite
aluminium carbonate
hydrochloric acid
sodium hydroxide
sulfuric acid
ammonium hydroxide
hydrogen
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Assignment 3
a) Using your Periodic Table, give an example for each of the following:
a noble gas
a metal with a valency of 1
a non-metal with a valency of 2
a liquid metal at room temperature
a liquid non-metal at room temperature
a non-metal which exists as diatomic molecules
a transition metal
a man-made metal
a radioactive metal
b) Find and record the atomic masses for the following elements
iron ……………….
iodine…………………… helium ……………………
lead ……………….
potassium ……………… sulfur ……………………..
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Assignment 4
a) Balance the following equations
HCl(aq) + Fe(s)
 FeCl2(aq) + H2(g)
(NH4)2Cr2O7(s)  Cr2O3(s) + N2(g) + H2O(g)
C4H10(g) + O2(g)
 CO2(g) +
H2O(l)
b) Describe how you would make up the following solutions for a senior chemistry class
500 mL of a 0.10 mol/L NaCl solution
500 mL of a 0.01 mol/L HCl solution from concentrated (10 M) HCl
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Assessment 2
Assignment 1
PART A
Given the following data and formulae, calculate the values of the unknown:
(1 mark each)
V 
1.
nRT
P
P = 1.22, R = 0.0821, T = 298, n = 0.0173
Answer
V= _________________
n 
2.
m
F
F = 125, n = 2.45 x 10-3
Answer
3.
m= _________________
E = 2  d (n + a)2
d = 39.4, n = 5.25 x 104, a = 1.67 x 10-3,  is pi (see your calculator)
Answer
E= _________________
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4.
C=
f [( V1 - V2) - (B1- B2)]
p
f = 0.1, p = 0.5, V1 = 9.7, V2 = 3.4, B1 = 1.2, B2 = 0.9
Answer
C= _________________
Write 0.000234 in scientific notation
Answer
= _________________
Write 4.56 x 104 as a decimal number
Answer
= _________________
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PART B
Perform the following conversions:
(1 mark each)
7.
789 nm to m
Answer
8.
4.25 x 10-4 kL to mL
Answer
9.
_________________ mL
270 mm2 to m2
Answer
10.
_________________ m
_________________ m2
62.3 kg to pounds
(1 pound = 454g)
Answer
_________________ pounds
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PART C
Plot a calibration graph on the supplied graph paper, given the data below:
(2 marks)
Concentration (mg/L)
Reading
0
0.01
2
0.18
4
0.35
6
0.55
8
0.74
Using your graph, determine the concentration of a sample with a reading of 0.44.
(1 mark)
Sample concentration
Determine the concentration of a sample, given the data below for a calibration graph.
(1 mark)
slope
25.4
y-intercept
1.5
sample reading
89.4
Sample concentration
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Assignment 2
PART A
1
How many moles are contained in 1.234
g of NaOH (FW = 40.0)?
2
What mass is 3.672 x 10-3 moles of lead
chloride
(2 marks each)
(FW = 278.1)?
3
Calculate the molarity of a solution
which has
3.84 x 10-4 moles in 25 mL of solution.
4
What molarity is produced when 1.329 g
of NaCl is dissolved in 500 mL of
solution?
5
What mass of NaCl (FW 58.44) would be
required to make 111 mL of 0.111 M
solution?
6
What is the concentration (in g/L) when
456 mg of KCl is dissolved in 1.5L of
solution?
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7
What mass of KNO3 would be required
to prepare 200 mL of 100 mg/L?
8
What is the concentration (in %w/w,
g/100g) of salt in soil if 0.0935 g of salt
were found in 22.763 g of soil?
9
What mass of lead is present in 350 g of
waste which has a lead concentration of
75 mg/kg?
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PART B
(2 marks each)
1.
What is the concentration of the diluted solution if 10 mL of 500
mg/L solution is diluted to 250 mL?
2.
What volume of 4 g/L solution is required to make 2 L of 0.05 g/L
solution?
3.
What is the concentration of a solution if it was diluted by 25 mL to
100 mL and the diluted solution had a concentration of 3.56 mg/L?
4.
To what volume should 10 mL of 0.05 M solution be diluted, so
that the diluted solution has a concentration of 0.0025 M?
5.
How many moles are present in 250 mL of solution if a 20 mL
aliquot of it contained 2.75 x 10-3 moles?
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