inorganic chemistry general and overall

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Inorganic Chemistry, tips and summary Section A. Constructing equations. Please realize that concepts learnt in other topics are still useful. The most important thing is to figure out the type of reaction. There are a few usual t ypes of reactions, namely 1. Acid-base 2. Redox 3. Precipitation 4. Decomposition 5. Hydrolysis 6. Displacement 1. Acid-base reactions. Look for H + donors or H + acceptors. Look for OH - donors or OH - acceptors. Eg. Metal oxides are basic: MO + 2HCl  ─ > MCl 2 + H 2 O The O 2- ion is actually a H + acceptor : O 2- + H 2 O ─ > 2OH -  Aluminium oxide and hydroxide are amphoteric and can act as acid. Al 2 O 3 + 2NaOH + 3H 2 O  ─ > 2Na[Al(OH) 4 ] Al(OH) 3 + :OH -   ─ > [Al(OH) 4 ]  ─  Lewis acid (accepts lone-pair from OH  ─ ) Notice how the two salts are the same… If you are constructing equation, focus on the ions that transfer H + or OH  ─  first. Balance charge with OH  ─  or H + or use counter-ions (such as Na) Always balance H or O with H 2 O LAST. 2. redox reactions: Watch out for i) common oxidizing agents (with particles in high oxidation s tate) eg MnO 4  ─ (Mn is +7) ii) common reducing agents ( metals, and particles at low oxidation state) eg Na, I  ─  iii) period trends - oxidizing power increase ac ross period (electronegativity increases) -reducing power decreases across period Group trends - Reducing power increase down a group (I  ─ is best reducing agt among halides, K is more reducing than Na in Gp I) - Oxidizing power increase up a group (F is best oxidizing agt among halogens) To construct equations : you can use the half equation mtd from notes (which is slow) or use this: i) Identify the starting material and product in the redox reaction. ii) Determine the in oxidation states, from there, determine the no. of e  ─ transferred. iii) Determine the stoichiometric ratio of the reagents by using the LCM of no of e  ─ transferred. iv) make sure the non H and O atoms are balanced first. v) balance charge with H + or OH  ─  

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Page 1: Inorganic Chemistry General and Overall

8/3/2019 Inorganic Chemistry General and Overall

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Inorganic Chemistry, tips and summary 

Section A.Constructing equations. Please realize that concepts learnt in other topics are still useful. The mostimportant thing is to figure out the type of reaction.

There are a few usual types of reactions, namely1. Acid-base2. Redox

3. Precipitation4. Decomposition5. Hydrolysis6. Displacement

1. Acid-base reactions. Look for H+ donors or H+ acceptors. Look for OH- donors or OH- acceptors.

Eg. Metal oxides are basic: MO + 2HCl  ─ > MCl2 + H2OThe O2- ion is actually a H+ acceptor : O2- + H2O  ─ > 2OH- 

Aluminium oxide and hydroxide are amphoteric and can act as acid.

Al2O3 + 2NaOH + 3H2O  ─ > 2Na[Al(OH)4]

Al(OH)3 + :OH-  ─ > [Al(OH)4] ─  

Lewis acid (accepts lone-pair from OH ─ )

Notice how the two salts are the same…

If you are constructing equation, focus on the ions that transfer H+ or OH ─ 

 first. Balance charge with OH ─  or H+ or use counter-ions (such as Na)

Always balance H or O with H2O LAST.

2. redox reactions: Watch out fori) common oxidizing agents (with particles in high oxidation state) eg MnO4

 ─  (Mn is +7)ii) common reducing agents ( metals, and particles at low oxidation state) eg Na, I ─  

iii) period trends- oxidizing power increase across period (electronegativity increases)-reducing power decreases across period

Group trends- Reducing power increase down a group (I ─  is best reducing agt among halides, K is more reducingthan Na in Gp I)- Oxidizing power increase up a group (F is best oxidizing agt among halogens)

To construct equations: you can use the half equation mtd from notes (which is slow) or use this:

i) Identify the starting material and product in the redox reaction.

ii) Determine the ∆ in oxidation states, from there, determine the no. of e ─  transferred.

iii) Determine the stoichiometric ratio of the reagents by using the LCM of no of e ─  transferred.

iv) make sure the non H and O atoms are balanced first.

v) balance charge with H+ or OH ─  

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vi) balance H or O atoms with H2O last.

This is extremely useful for GpVII and transition elts.

Example 2.1 Disproportionation of Cl2 in hot NaOH. (must remember oxidation state 0 ─ > -1, 0  ─ > +5)

i) Cl2  ─ > Cl ─  + ClO3 ─  

ii) Cl takes in 1 e ─  to give Cl ─ , Cl gives out 5 e ─  to become ClO3 ─  (V)

The mol ratio is 5:1 (LCM is 5)

iii) Cl2  ─ > 5Cl ─ + ClO3 ─  

iv) Balance non O, H atom first

3Cl2  ─ > 5Cl ─ + ClO3 ─  

v) 6 –ve charge on RHS, so, 3Cl2 + 6OH ─   ─ > 5Cl ─ + ClO3 ─  

vi) balance H and O with H2O: 3Cl2 + 6OH ─   ─ > 5Cl ─ + ClO3 ─ + 3H2O 

Example 2.2 Reaction of NaBr and NaI with conc. H2SO4 

Recognize that a) H2SO4 has acidic H+, 1st proton is very acidic, so acid base rxn takes place first.b) the sulphur in H2SO4 is +6, so it has some oxidizing ability.c) I ─  is a good reducing agent.

Acid-base: NaBr + H2SO4  ─ > NaHSO4 + HBr

Redox : HBr (-1) can reduce S from +6 to +4 (in SO2)

HBr + H2SO4  ─ > Br2 + SO2 (-1) (+6) (0) (+4) (1 e ─  for I, 2 e ─  for S, ratio should be 2:1)

2HBr + H2SO4  ─ > Br2 + SO2 

balance H or O with H2O 2HBr + H2SO4  ─ > Br2 + SO2 + 2H2O 

Acid-base: NaI + H2SO4  ─ > NaHSO4 + HI

Redox : HI(-1) is a good reducing agent and can even oxidize S from +6 to -2 (in H2S)

HI + H2SO4  ─ > I2 +H2S(-1) (+6) (0) (-2) (1 e ─  for I, 8 e ─  for S, ratio is 8:1)

8HI + H2SO4  ─ > 4I2 + H2S

Balance H or O with H2O 8HI + H2SO4  ─ 

> 4I2 + H2S + 4H2O 

You can see right now if you have a good ability to construct equations, memory work is reduced/madeeasier when studying inorganic chem. You should not simply memorize equations withoutunderstanding what sort of reactions are taking place, and how the starting materials become theproducts.

Predicting if a redox reaction takes place: Use Eo value.

For spontaneous reaction Eo >0, or ∆∆∆∆G<0

Note: ∆∆∆∆G = -nFE

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3. Precipitation

You have to be familiar with the solubilities of ionic cpds.

Soluble Sparingly soluble (Often known as “insoluble”)

All nitrates (NO3−) Nil.

Most halides (Cl−, Br−, I−) Exceptions: Halides of Ag+, Pb2+, Cu2+

Most sulphates (SO42−) Exceptions: Sulphates of Pb2+, Ca2+, Ba2+ 

Exceptions: Oxides of Na+, K+, NH4+ and

larger group II cations, eg. Ca2+, Sr2+, Ba2+.Most oxides (O2−) 

Exceptions: Hydroxides of Na+, K+ and largergroup II cations, eg. Ca2+, Sr2+, Ba2+.

Most hydroxides (OH−) 

Exceptions: Sulphides of Na+, K+, NH4+ and

group II cations, eg. Mg2+, Ca2+, Sr2+, Ba2+.Most sulphides (S2−)

Exceptions: Carbonates/ phosphates of Na+,K+ and NH4

+.Most carbonates (CO3

2−) & phosphates (PO43−) 

Exceptions: Chromates of Na

+

, K

+

, NH4

+

andMg2+ (Small group II cations) Most chromates (CrO4

2−

Spot some trends above. For example, all Gp I metal salts are soluble.

Be aware especially of hydroxides, since an acid-base problem can incorporate solubility.For example, if a reaction produced hydroxide, and metal ions are present, there is a high chance ofmetal precipitate forming.

For a good summary and examples: Refer to pg 11-12 of your Solubility Equilibria notes, and the QAsummary of your Transition element chemistry notes.

4. Decomposition 

Watch out for 1 compound reacting by itself, to break down to give more compounds/elements.Usually stable small molecules of gases are released. Stable solids with high LE such as oxides areusually residue.

You are expected to be familiar with the decomposition of Gp II nitrates, carbonates, hydroxides.

For unfamiliar decomposition reactions, clues are usually provided by the question. Read carefully anduse the context. Many decomposition reactions are also redox reactions.

Example 4.1Heating of ammonium dichromate causes decomposition to the oxide Cr2O3 and other gases. Constructthe equation for the reaction.

SolutionCr2O7

2- (Cr +6) Cr2O3 (Cr+3); Cr is reduced. In decomposition reaction, if one species is reduced,another species from the same cpd is oxidized. Each Cr gains 3 electrons.

The only other species left behind is the ammonium ion. Note also that since the mol ratio of NH4+ to

Cr2O72- is 2:1 [MF: (NH4)2Cr2O7] , each ammonium ion needs to lose 3 electrons.

Probable oxidation of NH4+ (N: -3) to N2 (N: 0)

(NH4)2Cr2O7  Cr2O3 + N2 + 4H2O(-3) (+6) (+3) (0)

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 5. Hydrolysis

Involves splitting water. 2 major types for inorganic chem..

1) For ionic cpds.Cation hydrolysis. Cation with high charge density M3+, have high polarizing power. Polarization ofwater molecules take place and the O-H bond will break more easily, releasing H+.

2) For covalent cpds. Eg Covalent chlorides.Similar 

to nucleophilic substitution (please do not use thisterm!, simply name the reaction as hydrolysis).

Water donates lone pair (as a nucleophile) to central atom eg, P in PCl3, and HCl is released.

For period 3 *low lying vacant d orbitals easily accept lone pair electrons from water, and hydrolysistakes place readily

Section B Learning inorganic chem. 

1. Don’t go straight into memorizing equations!

Key approaches:

-Big picture view first! Use general trends, relate structure + bonding to property, and familiarize withgeneral character of the compound/element.

-Relate specific reactions/equations to these big picture properties- It is more important to remember the identity of the starting material and product and type ofreaction; the equation can be constructed on the spot.

a. Structure (ionic vs covalent, oxidation state, Zeff) influences property (acidity, basicity, reducing power,oxidizing power)

2. trends.Eg. Oxides and chlorides across period, increasing covalency, increasing acidity.

-Ionic oxide, O2 ─ 

, basic. Covalent oxide, adds water to give XOOH, acidic-Ionic chlorides, Cl ─ doesn’t do much, Mx+ has increasing polarizing power, cation hydrolysis.

- Covalent chlorides, hydrolysis with water (like nucleophilic sub.) acidic solution

Eg. Gp II, polarizing power of M2+ decreases going down the group

Eg. Gp VII, oxidizing power of X2 is greatest at top of group, reducing power of X ─  is greatest at bottom ofgroup

3. When solving questions, make use of characteristic chemistry , as well as context and clues fromthe question.Eg.

1)Metal ions, lookout for : cation hydrolysis, precipitation, complex formation, redox

2) OH ─ , lookout for: neutralization, precipitation with metal ions, hydrolysis of organic/covalent cpds,redox reactions (sets alkaline conditions, affects equilibrium, especially for gp VII)

3) Gp VII, redox reactions.eg

X2 is oxidizingX2 can disproportionate in OH ─  X ─  is reducing.(reactions with H2SO4)

Precipitation. X ─  with Ag+, Pb2+