Transcript
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Qualitative Analytical ChemistryDr Mark Selby

E-Block E413D (GP)

[email protected]

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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

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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

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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

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Tools for writing reportsThe nuts and bolts of word processing for chemistry

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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/Resources/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.

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Office 365 University

Handy cloud storage . Always have your work with you.

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Wolfram Alpha

Wolfram Alpha solves chemical equations and provides comprehensive chemical data.

Apps for Android and iPad.

Handy for Mol. Wt.

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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

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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

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Things you should already knowBut probably need to revise!

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4. Oxyanions (polyatomic anions containing oxygen) are named with -iteor -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

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Monatomic anions:

Cl− chlorideS2− sulfideP3− phosphide

Polyatomic ions:

NH4+ ammonium

H3O+ hydronium NO3

− nitrate NO2

− nitrite ClO− hypochlorite ClO2

− chloriteClO3

− chlorate ClO4

− perchlorate SO3

2− sulfite SO4

2− sulfate HSO3

− hydrogen sulfite

HCO3− hydrogen carbonate

CO32− carbonate

PO43− phosphate

HPO42− hydrogen phosphate

H2PO4− dihydrogen phosphate

CrO42− chromate

Cr2O72− dichromate

BO33− borate

AsO43− arsenate

C2O42− oxalate

CN− cyanide SCN− thiocyanate MnO4

− permanganate

List of common ion names

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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-

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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 idewith 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

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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

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Acid Formula Anions Notes

Sulfuric acid H2SO4 sulfate ,SO42- and hydrogen

sulfate, HSO4-

Thiosulfuric acid H2S2O3 thiosulfate, S2O32− Aqueous solutions

decompose:{write equation}†

Sulfurous acid H2SO3

Sulfite, SO32- and hydrogen

sulfite, HSO3-

Aqueous solutionsdecompose:{write equation}†

Oxyacids of sulfur

Only the most common oxyacids are shown in the table below:

† Hint: see Prac. Manual pages 25 and 26

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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

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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

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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. Polydentateligands (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: NH3becomes ammine; H2O becomes aqua or aquo; CO becomes carbonyl; NO becomes nitrosyl.

Naming complexes

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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

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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

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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

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Safety in the Laboratory

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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

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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

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Remember the triple AAA:

Always add acids to water: never the other way around

Hazard warning symbols

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Lecture ContentQualitative Inorganic Analysis 1

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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

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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

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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)

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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

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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

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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

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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

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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

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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

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End of slides


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