The document outlines the curriculum and objectives for a qualitative analytical chemistry course at QUT, emphasizing the importance of chemical literacy and practical laboratory skills. It offers guidance on report writing, chemical nomenclature, and safe handling practices for acids and bases. Additionally, it includes information on various chemical compounds, their properties, and proper procedures for lab safety.

1. Qualitative Analytical Chemistry
Dr Mark Selby
E-Block E413D (GP)
m.selby@qut.edu.au

2.  Increasing chemical literacy
 Developing skills and knowledge that are relevant to
chemical industry and research.
 Achieving this:
 Practical skills through the laboratory program
 Developing chemical concepts (mental models) through
simulations --- a work in progress
 Alignment of fundamental principles taught in lectures
with practical work in the laboratory (where possible)
Objectives

3. Literacy, in its most common usage, is defined as the ability to
read and write. These are basic skills and the absence of one or
both is considered to be a handicap in an industrialized society.
Chemical Literacy

4.  See information on pages 12 – 17 of the Practical Manual and the
notes under the heading “Review” in these PowerPoint slides.
 Handwritten reports are acceptable and will not result in any
reduction in grades.
 Over the course of the semester in this unit you should be
developing your skills in writing reports that will meet a standard
acceptable to industry or research.
 Improvement in writing reports is to be judged by the structuring of
the report, the standard of written communication, quality of
recorded observations and discussion of results, and the use of
chemical equations.
 The mechanics of writing your report (i.e., handwritten versus word-
processing) is of secondary importance (but word processing may be
expected in Industry or research).
Writing chemical equations

5. Tools for writing reports
The nuts and bolts of word processing for chemistry

6. The RSC font:
 The Royal Society of Chemistry
font is a specialist chemistry font
that may be downloaded free and
used on PCs and Macs. This font
allows chemical symbols to be
introduced easily into Word
documents.
 http://www.rsc.org/Education/Teachers/Re
sources/Font.asp
Fonts for Chemistry
Chem97 Font:
 ChemFont97 is a Windows font package that simplifies the entry of chemical equations
and notation. The font includes all upper and lower case Greek characters, superscripts,
subscripts, many chemistry-specific symbols like reaction arrows. ChemFont97 comes in
two styles: serif, which is like Times New Roman, and sans-serif, which is like Arial.

7. Office 365 University
Handy cloud storage . Always
have your work with you.

8. Wolfram Alpha
Wolfram Alpha solves
chemical equations and
provides comprehensive
chemical data.
Apps for Android and
iPad.
Handy for Mol. Wt.

9.  Collaborate but don’t plagiarise. Make sure you understand
the difference!
 See for instance:
 http://www4.caes.hku.hk/writing_turbocharger/collaborating/
default_answers.htm
Collaborate but don’t Plagiarise

10. 1. Single atom anions are named with an -ide suffix: for
example, F− is fluoride.
2. Compounds with a positive ion (cation), the name of the
compound is simply the cation's name, followed by the
anion. For example, CaF2 is calcium fluoride.
3. Cations able to take on more than one positive charge
are labeled with Roman numerals in parentheses. For
example, Cu+ is copper(I), Cu2+ is copper(II).
 An older, out-dated notation is to append -ous or -ic to the
root of the Latin name to name ions with a lesser or greater
charge. Under this naming convention, Cu+ is cuprous and
Cu2+ is cupric.
IUPAC nomenclature

11. Things you should already know
But probably need to revise!

12. 4. Oxyanions (polyatomic anions containing oxygen) are named with -ite
or -ate, for a lesser or greater quantity of oxygen. For example, NO2
− is
nitrite, while NO3
− is nitrate. If four oxyanions are possible, the
prefixes hypo- and per- are used: hypochlorite is ClO−, perchlorate is
ClO4
−.
5. The prefix bi- is an out-dated way of indicating the presence of a single
hydrogen ion, as in "sodium bicarbonate" (NaHCO3). The preferred
method specifically names the hydrogen atom. Thus, NaHCO3 would
be called "sodium hydrogen carbonate".
6. The prefix thio indicates the substitution of oxygen by sulfur, so that
thiosulfate ion is a sulfate ion SO4
2- with one oxygen replaced by a
sulfur as in S2O3
2-.
 The preferred IUPAC name for the protonated species H2S2O3 is thiosulfuric
acid.
IUPAC nomenclature

13. Monatomic anions:
Cl− chloride
S2− sulfide
P3− phosphide
Polyatomic ions:
NH4
+ ammonium
H3O+ hydronium
NO3
− nitrate
NO2
− nitrite
ClO− hypochlorite
ClO2
− chlorite
ClO3
− chlorate
ClO4
− perchlorate
SO3
2− sulfite
SO4
2− sulfate
HSO3
− hydrogen sulfite
HCO3
− hydrogen carbonate
CO3
2− carbonate
PO4
3− phosphate
HPO4
2− hydrogen phosphate
H2PO4
− dihydrogen phosphate
CrO4
2− chromate
Cr2O7
2− dichromate
BO3
3− borate
AsO4
3− arsenate
C2O4
2− oxalate
CN− cyanide
SCN− thiocyanate
MnO4
− permanganate
List of common ion names

14.  Hydrates are ionic compounds that have absorbed water. They
are named as the ionic compound followed by a numerical prefix
and -hydrate. The numerical prefixes used are listed below:
 For example, CuSO4 · 5H2O is "copper(II) sulfate pentahydrate".
Naming hydrates
mono-
di-
tri-
tetra-
penta-
hexa-
hepta-
octa-
nona-
deca-

15.  Acids are named by the anion they form when dissolved in water. If
an acid forms an anion ending in ide, then its name is formed by
adding the prefix hydro to the anion's name and replacing the ide
with ic. Finally the word acid is appended.
 For example, hydrochloric acid forms a chloride anion. With sulfur,
however, the whole word is kept instead of the root: i.e.:
hydrosulfuric acid.
 Secondly, anions with an -ate suffix are formed when acids with an -
ic suffix are dissolved, e.g. chloric acid (HClO3) dissociates into
chlorate anions to form salts such as sodium chlorate (NaClO3);
 anions with an -ite suffix are formed when acids with an -ous suffix
are dissolved in water, e.g. chlorous acid (HClO2) disassociates into
chlorite anions to form salts such as sodium chlorite (NaClO2).
Naming acids

16.  The four oxyacids of chlorine are called hypochlorous acid
(HClO), chlorous acid (HClO2), chloric acid (HClO3) and
perchloric acid (HClO4).
 Their respective conjugate bases are the hypochlorite
(ClO-), chlorite (ClO2
-), chlorate (ClO3
-) and perchlorate
(ClO4
-) ions.
 The corresponding potassium salts are potassium
hypochlorite (KClO), potassium chlorite (KClO2), potassium
chlorate (KClO3) and potassium perchlorate (KClO4).
Oxyacids of chlorine

17. Acid Formula Anions Notes
Sulfuric acid H2SO4 sulfate ,SO4
2- and hydrogen
sulfate, HSO4
-
Thiosulfuric acid H2S2O3 thiosulfate, S2O3
2− Aqueous solutions
decompose:
{write equation}†
Sulfurous acid H2SO3
Sulfite, SO3
2- and hydrogen
sulfite, HSO3
-
Aqueous solutions
decompose:
{write equation}†
Oxyacids of sulfur
Only the most common oxyacids are shown in the table below:
† Hint: see Prac. Manual pages 25 and 26

18.  When the metal has more than one possible ionic charge or
oxidation number the name becomes ambiguous.
 In these cases the oxidation number (the same as the charge) of
the metal ion is represented by a Roman numeral in parentheses
immediately following the metal ion name.
 For example in uranium(VI) fluoride the oxidation number of
uranium is 6. Another example is the iron oxides. FeO is iron(II)
oxide and Fe2O3 is iron(III) oxide.
 An older system used prefixes and suffixes to indicate the
oxidation number, according to the following scheme (next
page):
Traditional naming

19. This system has partially fallen out of use, but survives in the
common names of many chemical compounds: e.g., "ferric chloride"
(instead calling it "iron(III) chloride") and "potassium
permanganate" (instead of "potassium manganate(VII)").
Traditional naming
Oxidation state Cations and acids Anions
Lowest hypo- -ous hypo- -ite
-ous -ite
-ic -ate
per- -ic per- -ate
Highest hyper- -ic hyper- -ate

20.  When naming a complex ion, the ligands are named before the
metal ion.
 Write the names of the ligands in alphabetical order (numerical
prefixes do not affect the order.)
 Multiple occurring monodentate ligands receive a prefix according to the
number of occurrences: di-, tri-, tetra-, penta-, or hexa. Polydentate
ligands (e.g., ethylenediamine, oxalate) receive bis-, tris-, tetrakis-, etc.
 Anions end in ido. This replaces the final 'e' when the anion ends with
'-ate', e.g. sulfate becomes sulfato. It replaces 'ide': cyanide becomes
cyanido.
 Neutral ligands are given their usual name, with some exceptions: NH3
becomes ammine; H2O becomes aqua or aquo; CO becomes carbonyl; NO
becomes nitrosyl.
Naming complexes

21.  Write the name of the central atom/ion. If the complex is an
anion, the central atom's name will end in -ate, and its Latin
name will be used if available (except for mercury).
 If the central atom's oxidation state needs to be specified, write
it as a Roman numeral in parentheses.
 Name cation then anion as separate words.
Examples:
[NiCl4]2− tetrachloridonickelate(II) ion
[CuNH3Cl5]3− amminepentachloridocuprate(II) ion
[Cd(en)2(CN)2] dicyanidobis(ethylenediamine)cadmium(II)
[Co(NH3)5Cl]SO4 pentaamminechloridocobalt(III) sulfate
Naming complexes

22. Soluble Insoluble
Group I and NH4
+ compounds
Carbonates (Except Group I, NH4
+ and
uranyl compounds)
Nitrates, chlorates (all are highly soluble)
Sulfites (Except Group I and NH4
+
compounds)
Acetates (Ethanoates) (Except Ag+
compounds)
Phosphates (Except Group I and NH4
+
compounds)
Chlorides, bromides and iodides (Except
Ag+, Pb2+, Cu+ and Hg2
2+)
Hydroxides and oxides (Except Group I,
NH4
+, Ba2+, Sr2+ and Tl+)
Sulfates (Except Ag+, Pb2+, Ba2+, Sr2+ and
Ca2+)
Sulfides (Except Group I, Group II and
NH4
+ compounds)
Chromates (Except Na2CrO4, K2CrO4,
(NH4)2CrO4, and MgCrO4)
Solubility Rules

23. Half reaction E° (in volt) Half reaction E° (in volt)
Li+
(aq) + e– ∏ Li(s) –3.02 Co2+
(aq) + 2e– ∏ Co(s) –0.28
K+
(aq) + e– ∏ K(s) –2.93 Ni2+
(aq) + 2e– ∏ Ni(s) –0.23
Ca2+
(aq) + 2e– ∏ Ca(s) –2.87 Sn2+
(aq) + 2e– ∏ Sn(s) –0.14
Na+
(aq) + e– ∏ Na(s) –2.71 Pb2+
(aq) + 2e– ∏ Pb(s) –0.13
Mg2+
(aq) + 2e– ∏ Mg(s) –2.34 2H+
(aq) + 2e– ∏ H2(g) 0.00
Al3+
(aq) + 3e– ∏ Al(s) –1.67 Cu2+
(aq) + 2e– ∏ Cu(s) +0.34
Mn2+
(aq) + 2e– ∏ Mn(s) –1.03 Ag+
(aq) + e– ∏ Ag(s) +0.80
Zn2+
(aq) + 2e– ∏ Zn(s) –0.76 Au+
(aq) + e– ∏ Au(s) +1.68
Fe2+
(aq) + 2e– ∏ Fe(s) –0.44
The electrochemical series

24. Safety in the Laboratory

25. Safe Handling of Acids and Bases
 There are a number of proper procedures for the safe handling of acids and bases
that you need to know because you’ll be working with them quite a bit. (These
guidelines are also listed in your Practical Manual.)
 Both acids and bases can be corrosive to human tissue. When concentrated, they can
react with tissue and break it down. In general, the more concentrated the acid or
base happens to be, the more hazardous it is. Although the more concentrated acids
and bases are the most dangerous ones, don't ignore the dilute ones.
 You must be particularly careful about getting them in your eyes. It is compulsory to
wear safety glasses when handling either acids or bases. But if you do get any in your
eyes, let the demonstrator know and flush it out immediately with lots of water,
several minutes worth. There are eye washes in the lab. You will learn where they are,
and how to use them, in the online induction.
Precautions

26. Safe Handling of Acids and Bases
 Suppose you get some acid or base on you, other than in your eyes. The procedure is
the essentially the same: flush that area immediately for several minutes with water
and consult the demonstrator for further advice. If you should be unfortunate
enough to spill it all over you, use the safety shower in the lab.
 If you spill acid or base on the lab bench top or on the floor, treat it immediately. If it
is an acid, first neutralize it with sufficient sodium hydrogen carbonate, (commonly
known as baking soda). We have spill kits for use with extensive spills available in the
lab or in the adjoining prep room (check with demonstrator or technical staff).
 The quantities and concentrations of acids and bases used in the exercises at QUT are
permissible to flush down the drain. However, the quantity involved in a spill should
be neutralized before disposal (see technical staff).
Precautions

27.  Remember the triple AAA:
 Always add acids to water: never the other way around
Hazard warning symbols

28. Lecture Content
Qualitative Inorganic Analysis 1

29.  Also known as inorganic acids or mineral acids and include: Hydrochloric
acid: Nitric acid (dilute), Phosphoric acid, Sulfuric acid (dilute), Boric acid,
Hydrofluoric acid and Hydrobromic acid.
 Inorganic acids are generally soluble in water with the release of hydrogen
ions. The resulting solutions have pHs of less than 7.0.
 Acids neutralize chemical bases (for example: amines and inorganic
hydroxides) to form salts. Neutralization occurs as the base accepts
hydrogen ions that the acid donates.
 Neutralization can generate dangerously large amounts of heat in small
spaces. The dissolution of acids in water or the dilution of their
concentrated solutions with additional water may generate significant heat;
the addition of water often generates sufficient heat in the small region of
mixing to cause some of the water to boil explosively. The resulting
"bumping" spatters the acid.
 These materials react with active metals, including such structural metals as
aluminum and iron, to release hydrogen, a flammable gas.
Non-oxidising acids

30. Test for Test method Observations Equations
H2(g) lighted splint “pop” {write equation}
O2(g) Glowing splint Re-ignites {write equation}
CO2(g) limewater White ppt {write equation}
HCl(g) Moist blue litmus paper Turns red {write equation}
NH3(g) Moist red litmus paper Turns blue {write equation}
SO2(g) Moist dichromate paper Turns green {write equation}
Cl2(g) Pungent gas - Drop of AgNO3 White ppt {write equation}
Br2(g) Orange-brown gas - Drop of AgNO3 Cream ppt {write equation}
NO2(g) Orange-brown gas No simple test {write equation}
NO(g) Colourless → orange brown gas Colourless in the
absence of air
{write equation}
H2S(g) moist lead ethanoate (acetate) paper Turns black {write equation}
Qualitative tests for gases

31.  A basic oxide is an oxide that shows basic properties (in
opposition to acidic oxides) and that either:
 reacts with water to form a base; or
 reacts with an acid to form a salt.
 Examples include:
 Sodium oxide which reacts with water to produce sodium hydroxide
 Magnesium oxide, which reacts with hydrochloric acid to form
magnesium chloride
 Copper oxide, which reacts with nitric acid to form copper nitrate
 Basic oxides are oxides mostly of metals, especially alkali and alkaline
earth (Group I and II) metals.
Basic oxides (and hydroxides)

32.  An amphoteric substance is a compound that can react as an acid as
well as a base. The word is derived from the Greek word amphoteroi
(ἀμφότεροι) meaning "both".
 Many metals (such as zinc, tin, lead, aluminium, and beryllium) and most
metalloids have amphoteric oxides or hydroxides. Amphoteric
substances can either donate or accept a proton.
 Zinc oxide (ZnO) reacts differently depending on the pH of the solution:
 In acids: ZnO + 2H+ → {write products}
 In bases: ZnO + H2O + 2OH- → {write products}
 This effect can be used to separate different cations, such as zinc from
manganese (next week: separation schemes).
Amphoteric oxides and hydroxides

33.  Aluminium hydroxide is as well:
 Base (neutralizing an acid): Al(OH)3 + 3HCl → {write products}
 Acid (neutralizing a base): Al(OH)3 + NaOH → {write products}
 Some other examples include:
 Aluminium oxide
 with acid: Al2O3 + 3H2O + 6H3O+(aq) →) {write products}
 with base: Al2O3 + 3H2O + 2OH-(aq) → 2 {write products}
 Lead oxide
 with acid: PbO + 2HCl → {write products}
 with base: PbO + Ca(OH)2 +H2O → {write products}
 Question: see if you can write ionic chemical equations that illustrate the
amphoteric nature of Beryllium hydroxide?
Amphoteric oxides and hydroxides

34.  An oxidizing acid is a Brønsted acid that is also a strong oxidizing agent.
All Brønsted acids can act as moderately strong oxidizing agents, because
the acidic proton can be reduced to hydrogen gas.
 However, some acids contain other structures that act as stronger
oxidizing agents than hydrogen. Generally, they contain oxygen in the
anionic structure.
 These include nitric acid, perchloric acid, chloric acid, chromic acid, and
concentrated sulfuric acid, among others.
 For example, copper metal cannot be oxidized by and dissolved in a non-
oxidizing acid, because it is lower on the reactivity series than acidic
hydrogen. However, an oxidizing acid such as nitric acid can oxidize the
copper, and allow it to dissolve.
Oxidizing acids

35.  Strongly electropositive metals, such as magnesium react with nitric acid as with other acids,
reducing the hydrogen ion.
 Mg + 2 H+ → {write products}
 With less electropositive metals the products depend on temperature and the acid concentration.
For example, copper reacts with dilute nitric acid at ambient temperatures with a 3:8 stoichiometry.
 3Cu + 8 HNO3 → {write products}
 The nitric oxide produced may react with atmospheric oxygen to give nitrogen dioxide.
 With more concentrated nitric acid, nitrogen dioxide is produced directly in a reaction with 1:4
stoichiometry.
 Cu + 4H+ + 2 NO3
− → {write products}
Passivation
 Although chromium (Cr), iron (Fe) and aluminium (Al) readily dissolve in dilute nitric acid, the
concentrated acid forms a metal oxide layer that protects the metal from further oxidation, which is
called passivation.
Nitric Acid

36.  Sulfuric acid reacts with most bases to give the corresponding sulfate. For example, the blue copper
salt copper(II) sulfate is prepared by the reaction of copper(II) oxide with sulfuric acid:
CuO(s) + H2SO4 (aq) → {write products}
 Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium
acetate, for example, displaces acetic acid, CH3COOH, and forms sodium hydrogen sulfate:
H2SO4 + CH3COONa → {write products}
 Similarly, reacting sulfuric acid with potassium nitrate can be used to produce nitric acid and
potassium hydrogen sulfate.
 Concentrated sulfuric acid reacts with sodium chloride, and gives hydrogen chloride gas and sodium
hydrogen sulfate:
NaCl + H2SO4 → {write products}
 Similarly with other halide salts
Sulfuric Acid

37.  Sulfuric acid reacts with most metals via a single displacement reaction to produce hydrogen gas
and the metal sulfate. Dilute H2SO4 attacks iron, aluminium, zinc, manganese, magnesium and nickel,
 but reactions with tin and copper require the acid to be hot and concentrated.
 Lead and tungsten, however, are resistant to sulfuric acid. The reaction with iron shown below is
typical for most of these metals, but the reaction with tin produces sulfur dioxide rather than
hydrogen.
Fe(s) + H2SO4(aq) → {write products}
Sn (s) + 2H2SO4(aq) → {write products}
 These reactions may be taken as typical: the hot concentrated acid generally acts as an oxidizing
agent whereas the dilute acid acts as a typical acid. Hence hot concentrated acid reacts with tin,
zinc and copper to produce the salt, water and sulfur dioxide, whereas the dilute acid reacts with
metals high in the reactivity series (such as Zn) to produce a salt and hydrogen.
Sulfuric Acid

38. End of slides