chemistry for first year university

7/21/2019 Chemistry for First Year University

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Chemistry Notes: Essential Knowledge forUniversity Chemistry of Any Level

By Josh Walker

Includes: Acids and Bases Bonding Electro Chemistry Equilibrium Rate of Reaction Redox Titrations

By Josh Walker


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Contents

Acids and Bases3

Bonding21

Electro Chemistry50

Equilibrium76

Rate of Reaction76

Redox97

Titrations105

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Acids and Bases

Introduction/Revision:

Strong Acids:And acid is generally defined as:

A substance that produces hydrogen ions or protons in water Arrhenius Theory A SUBSTANCE THAT DONATES PROTONS DURING A CHEMICAL REACTION

Bronsted/Lowrey theory.

A strong acid is a substance that completely ionises when dissolved in water:

If we look at HCl this is represented by:

However, the actual reaction is that of the proton reacting with the water molecule as below:

Some important strong acids are: HCl hydrochloric acid HNO3 Nitric Acid HClO4 perchloric acid HI iodic acid HBr bromic acid H2SO4 sulfuric acid.

It is important to note that when a strong acid ionises the concentration of the molecules is zero

you can only have a concentration of the ions.

We can categorize acids by the number of protons H+they produce when they ionise: Monoprotic ionise to produce one proton.

o HCl.

o HNO3 Diprotic ionise to produce two protons.

o H2SO4.

o H2C2O4.

Triprotic ionise to produce three protons.

o H3PO4.

When we show the ionisation of a diprotic or triprotic acid it is important we show it in steps.I.e.

the ionisation of sulfuric acid can be shown:

IT is important to note that for every mole of H2SO4you get around 1.6 moles of H+ion. So the

correspondence IS NOT A ONE TO ONE RATIO AS THE H2SUGGESTS.

HCl(aq)Water

! "! ! H+(aq) +Cl#

(aq)

HCl(aq) +H2O(l) !H3O+

(aq) +Cl"

(aq)

Hydronium Ion

First Ionisation: H2SO4(aq) !H

+

(aq) + HSO4"

(aq)

Second Ionisation: HSO4"

(aq)! H+

(aq) + SO4"

(aq)

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We also need to remember that because water molecules are polar they react with the H+ion to

form hydronium ions as shown below.

This is the bases for acids and strong acids.

Strong Bases:You can generally define bases or ALKALIS as:

A substance that produces hydroxide ions in solution Arrhenius Theory A SUBSTANCE THAT ACCEPTS PROTONS DURING A CHEMICAL REACTION

Bronsted/Lowrey theory.

A strong base like a strong acid is a substance that completely ionises into ions

As you can see NaOH is a strong base:

Basically ALL METAL HYDROXIDES ARE STRONG BASES.Even though they are not very

soluble they are still strong bases because WHEN they do dissolve they completely ionise.

Weak Acids:

Weak acids are simply the opposite from strong acids THEY ONLY PARTIALLY IONISE IN

SOLUTION.

Common weak acids that we need to remember are: ALL organic acids (CH3COOH, HCOOH) Carbonic Acid H2CO3 Sulfurous acid H2SO3 Hydrofluoric acid HF

If we look at acetic acid dissolving in 1.0l of water:

As you can see the concentration of the ions in the solution is far less than the concentration of the

molecules in the solution. SO THE ACID ONLY PARTIALLY IONISE, there is only a small

number of H+ions, which CAN be donated in a reaction and we have a weak base.

It is important to know that you cannot calculate the above concentration in the exam you can only

calculate them when given tables of their values.

Weak Bases:

As with weak acids we bases TEND TO REMAINS AS MOLECULES THEY ONLY

PARTIALLY IONISE.

NH3Is a weak base as it only partially ionises by accepting a proton from water:

METAL OXIDES also react or HYDROLYSE with water and are thus weak bases:

H2O(l) +H+

(aq)! H3O+

(aq)

NaOH(s)Water

! "!! Na+(aq) +OH#

(aq)

CH3COOH(l)! CH3COO!

(aq) +H+

(aq)

Molecules Ions Ions

0.99molL!1

0.01molL!1

0.01molL!1

NH3(g) +HOH(l)! NH4+

(aq) +OH!

(aq)

MgO(s) +H2O(l)! Mg2+

(aq) +2OH!

(aq)

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It is very important to note that almost ALL metal oxides are insoluble in water.So most of them

will not ionise at all.

Concentration Note:

It is important that you understand the difference between strong and weakand concentrated and

dilute.For example:

The hydrochloric acid in our stomachs is a very dilute STRONG acid. The tannic acid in a strong cup of tea is a concentrated WEAK acid.

So whether an acid is strong or weak is determined by how it ionises.

Whether an acid is concentrated or dilute determines how much of that acid is in the solution you

are using.

This is a common mistake in answering simple questions so make sure it is not one of yours.

Water Acting As Acid/Base/Neutral:

Water is a very special substance in that it can react as an acid, as a base or as a neutral substance.

Below are examples of it reacting in different ways:

Water as an Acid:

The most common example of this is water reacting with ammonia:

As you can see the WATER has DONATED a proton to the ammonia so it ACTS AS AN ACID.

Water as a Base:

Here a common example is water reacting with hydrochloric acid:

As you can see the WATER has ACCEPTED a proton from the HCl so it ACTS AS A BASE.

Water acting Neutral:

The standard equation for this is:

In this equilibrium equation the concentrations of H+and OH-are equal.

It is worth having these on your calculator as the WILL come in useful.

Non-Electrolytes:These are substance that dissolve in water but remain as molecules. So they DO NOT

DISSOCIATE OR IONISE.

Substances like this include: Glucose C6H12O6 Sucrose C12H22O11 Ethanol CH3CH2OH

With non-electrolytes you need to remember they do not conduct an electric current and are neither

acidic nor basic.

NH3(g) +H2O(l)! NH4+

(aq) +OH!

(aq)

HCl(aq)

+H2

O(l)! H

3

O+(aq)

+Cl!(aq)

H2O(l)! H+

(aq) +OH!

(aq)

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pH:

Introduction to pH:

pH is essentially looking at the concentration of the hydrogen ion in the solution and it is given by the

formula:

If we look at water dissociating into a small number of ions:

We know that there is a large number of water molecules but a very small amount of ions in fact:

For 1 MOLE of H2O molecules there are:

1 x 10-7moles of H+ions and 1 x 10-7moles of OH-ions.

So from this we can say that:

This is a very important point to remember because we KNOW THAT WATER IS NEUTRAL.

So Therefore it follows that THE CONCENTRATION OF H+IONS MULTIPLIED BY OH-

IONS MUST ALWAYS EQUAL 1.0 X 10-14.

It is far easier however to simply say:

Where pOH is given by .

Before we move on to look at calculating pH it is important to remember what the pH scale means:

NOTE: pH can be negative if the H+concentration is GREATER than 1mlL-1.

Worked examples of finding pH or concentration:

Below are a few examples including descriptions of what is being done on how to find the pH of

different solutions and the concentration of H+ions from the solution:

Example 1 The easiest:

Calculate the pH of 1.00 x 10-3molL-1HNO3:

WRITE IONIC EQUATION

So it is a one to one correspondence from molecules to ions so we do not need to alter given

concentration.

SUB INTO FORMULA

pH =! log H+

"# $%

H2O! H+

(aq) +OH!

(aq)

H+

!" #$ OH

%

!" #$= 1.0

&10

%14

pH + pOH = 14! log OH!"# $%

HNO3! H

+

+OH"

pH =! log 1.0"10!3#$ %& =3.0

By Josh Walker

0Higher concentration of

H+Acidic

7

Neutral

14Lower concentration of H

+Basic (alkaline)


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Example 2 Diprotic:

Calculate the pH of 0.040 molL-1H2SO4:

WRITE IONIC EQUATION

So we need to multiply concentration given by 2 to get concentration of H+

ionsthus it is 0.08molL-1

USING ADJUSTED CONCENTRATION SUB INTO FORMULA

Example 3 Simple Concentration:

What is the concentration of a pH 2.5 HClO4solution:

WRITE IONIC EQUATION TO CHECK IT IS 1 1 CORRESPONDENCE:

1 mole of the acid goes to 1 mole of the hydrogen ion so this is no problem.

SIMPLY WRITE AS 10-pH:

In this case this is

But this is not correct scientific notation

PLUG INTO CALCULATOR TO GET CORRECT FORM:

Example 4 pH from pOH:

Calculate the pH of 0.025 molL-1Ba(OH)2:

WRITE IONIC EQUATION TO CHECK CORRESPONDENCE:

1 mole of the reactant goes to two moles of OH -so we must alter the concentration of OH -to equal =

0.05molL-1.

CALCULATE pOH AS GIVEN ABOVE USING ADJUSTED CONCENTRATION:

USE pH=14-pOH TO FIND pH:

Example 5 Concentrations from pH:

A solution has a pH of 4.5. Find the concentration of H+and OH-separately:

WORK OUT pOH:

TREAT BOTH pH and pOH AS IN EXAMPLE THREE:

H2SO

4! 2H

+

+OH"

pH=! log 0.08[ ]= 1.1

HClO4! H

+

+ClO4

"

10!2.5

pH = 3.2 !10"3

Ba OH( )2! Ba

2+ +2OH"

pOH=! log 0.05[ ]= 1.3

pH = 14! 1.3 = 12.7

pOH = 14! 4.5 =9.5

H+!" #$ = 10

%4.5= 3.2&10%5molL%1

OH%

!" #$= 10

%9.5= 3.2

&10

%10molL

%1

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Example 6 Doing some Calculating:

5.0g of NaOH is dissolved in 2.5L of water. What is the pH:

FIND HOW MANY MOLES OF OH-THIS IS:

CALCULATE CONCENTRATION:

FIND pOH:

FROM THIS FIND pH:

It is essential that all the above worked examples are understood before we move on to the more

complicated topic of limiting reagent problems.

pH and Limiting Reagents:

In this course we only look at neutralisation reactions. Basically in a BADLY DONE

neutralisation reaction one of the reactants will not be all used up it will be in excess (XS) or

(INXS): If an ACID (H+) is IN EXCESS then the pH will be lower than 7. If a BASE (OH-) is IN EXCESS then the pH will be higher than 7.

Limiting reagent calculations are actually overly simple but below are two examples of the method

used:

Example 1 The simplest type:

100mL of 0.07mol L-1LiOH solution is mixed with 200ml of 0.02mol L-1HBr solution what is the

resultant pH:

FIRST IDENTIFY LIMITING REAGENT

CALCULATE REAMING MOLES OF SUBSTANCE INXS

Because OH-is in XS pH will be above 7.

CALCULATE CONCENTRATION OF THIS SUBSTANCE

You must here be aware that THE VOLUME HAS CHANGED!!!!!

HENCE FIND pH:

n NaOH( ) = 5

40=0.125mol

NaOH! Na+ + OH"

#n(OH" )=0.125mol

c=n

v=

0.125

2.5= 0.05molL

!1

pOH=! log 0.05[ ]= 1.3

pH = 14! 1.3 = 12.7

n(LiOH) =0.1!0.07 =0.007

LiOH" Li+

+OH#

$n(OH# ) =0.007mol

n(HBr) =0.2 !0.02 =0.004

HBr" H+

+ Br#

$n(H+ ) =0.004mol

EquationH+ +OH!"H2O

So because it is 1-1 H+ is the Limiting Reagent

n(OH!

) =0.007 !0.004

"n(OH!

) =0.003mol

c=

n

v=

0.003

0.1+0.2=

0.003

0.3=0.01molL

!1

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Example 2 The slightly more complicated type:

Assuming all species are soluble what is the final pH when 2.0L of 0.025mol L-1of Aluminium

hydroxide are added to 0.5L of 0.03 mol L-1oxalic solution:

IDENTIFY LIMITING REAGENT (Remember to multiply to find concentration of ions)

CALCULATE REMAINING MOLES OF SUBSTANCE IN EXCESS

Because OH-is in XS pH will be above 7.

CALCULATE THE CONCENTRATION OF THIS SUBSTANCE

HENCE FIND pH:

This is pretty much as complicated as equations involving limiting reagents and pH will get!!

Advanced pH problems:

Below are a few examples of interesting and challenging problems involving pH they are good to

know because they are common occurrences in exams:

Example 1 Mixing and Finding New pH:

Equal volumes of solutions pH=1.0 and pH=2.0 are mixed what is the final pH:

CALCULATE CONCENTRATION OF H+IN EACH SOLUTION ORIGINALLY

ASSUME VOLUME TO BE 1 AND THUS FIND MOLES OF H+IN EACH SOLUTION

FIND TOTAL NUMBER OF MOLES OF H+WHEN COMBINED

DIVIDE BY NEW VOLUME (2) TO SOLVE FOR FINAL CONCENTRATION

pOH= ! log 0.01[ ]=2" pH= 14 !2= 12.0

n(Al OH( )3)=2 !0.025= 0.05

Al OH( )3" Na

3+ +3OH#

$n(OH# )= 3 !0.05= 0.15mol

n(H2C2O4 ) =0.5 !0.03 =0.015

H2C2O4 "2H+

+C2O42#

$n(H+ ) =2 !0.015 =0.03mol

EquationH+ +OH!"H2O

So because it is 1-1 H+is the Limiting Reagent

n(OH!

) =0.15 !0.0.03

"n(OH!

) =0.12mol

c=n

v=

0.12

2 +0.5=

0.12

2.5=0.048molL

!1

pOH= ! log 0.048[ ]= 1.3" pH= 14 !1.3= 12.7

Solution.1= 10

!1

= 1"10!1

molL!1

Solution.2 = 10!2

= 1"10!2molL

!1

n=cv so

Solution.1 = 1!10"1mol

Solution.2 = 1!10"2mol

n(total) = 1!10"1 +1!10"2 =0.11mol

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of12 117 OXYACIDS (containing oxygen) ENGING IN IC:

o PRODUCE SALTS ENDING IN ATE

! Sulfuric acid produces sulfate salts.

! Phosphoric acid produces phosphate salts.

OXYACIDS ENDING IN OUS:

o PRODUCE SALTS ENDING IN ITE! Nitrous acid produces nitric salts.

It is vital that you understand and remember these rules.

Reaction of Acids And Bases:

We have done this section to death over the past few years but we need to look at it just one more time: An acid is a substance that produces hydrogen ions.

o Therefore it is only the H+ion that reacts in acid reactions.

Like before lets for one last time look at the reactions of acids and bases in their various forms with

explanations and all that jazz:

REACTIONS OF ACIDS AND BASES NEUTRALISATION

If the acid and base are in solution then not all ions will react so only those reacting need to be

included and you get the standard reaction.

When the acid reacts with the solid base however you need to include the entire basic compound

as a solid:

When we have a weak acid which only partially ionises you need to use the molecular form of

that acid i.e:

You need to know METAL OXIDES ARE SOLID BASES, these may hydrolyse with water to

form OH-ions, however all of them are solids so you get reactions like:

Those are all the neutralisation reactions you need to know.

REACTION OF ACIDS WITH METALS

All metals EXCEPT MERCURYare solid. When and IF CHECK E0TABLE a acid reacts

with a metal then the metal ions are dissolved and H2gas is produced:

IT IS WORTH NOTING THAT THIS IS ACTUALLY A REDOX REACTION. Because it

involves the transfer of an electron and not a H+transfer.

REACTION OF ACIDS WITH CARBONATES

Once again if the substances are in solution then the non reacting ions do not need to be includedso you get the standard reaction:

H+(aq) +OH!

(aq) "H2O(l)

H+

(aq)

+NaOH(s)

! Na+

(aq)

+H2

O(l)

CH3COOH(aq) +OH!

(aq) "CH3COO!

(aq) +H2O(l)

2H+(aq) +CaO(s)!Ca2+

(aq) +H2O(l)

2H+(aq) +Zn(s)! Zn2+

(aq) +H2(g)

2H+(aq) +CO32!

(aq) "CO2g) +H2O(l)

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If however the carbonate is a solid then you need to write it as a solid:

This does involve H+transfer so it is an acid reaction.

REACTION OF ACIDS WITH HYDROGEN CARBONATES

These reactions are exactly the same as those for carbonates including the part about writing asa solid. If however both substances are in solution then you get the standard reaction:

Make sure you note the difference between number f H+ions reaction.

AMMONIA AS A BASE

We know ammonia is a base so it is important we know how to show the reaction where it

ACCEPTS A PROTON:

Knowing this is essential to your acid and base theory but unlike previous years, this year it is

only the beginning.

Amphoteric Substances:

Amphoteric substances are substance that react with acids and bases.These substances are: Aluminium. Zinc. Chromium.

o AND THEIR HYDROXIDES AND OXIDES.

The acid reactions are just applications of those above but it is worthwhile knowing the reactionswith their BASE and putting them on the calculator so lets look at them now:

ALUMINIUM

Aluminium Metal

Aluminium Hydroxide

Aluminium Oxide

ZINC

Zinc Metal

Zinc Hydroxide

Zinc Oxide

2H+(aq) +K2CO3(s)! 2K+

(aq) +CO2g) +H2O(l)

H+(aq) +HCO3!

(aq) "CO2g) +H2O(l)

H+(aq) +NH3(g)! NH4+

(aq)

2Al(s)+2OH!

(aq)+6H2O(l) "2Al OH( )4!

(aq)+ 3H2( g)

Al OH( )3(s)

+ OH!(aq) "Al OH( )4!

(aq)

Al2O3(s)+2OH!

(aq)+ 3H2O(l) "2Al OH( )4!

(aq)

Zn(s)+2OH!

(aq)+2H2O(l) "Zn OH( )42!

(aq)+ H2( g)

Zn OH( )2( s)+2OH!

(aq) "Zn OH( )42!

(aq)

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CHROMIUM

Chromium Metal

Chromium Hydroxide

Chromium Oxide

These are all the amphoteric reactions you need to know.

Acidic, Basic and Neutral Salts:

Basic Theory:

The bases to understanding this topic lies in aspects we have just looked at in this set of notes: Salts are ionic compounds that dissociate into ions when they dissolve in water they are

all strong electrolytes. Acids and Bases may be strong electrolytes ionise completely in water. Acids or bases may also be weak electrolytes meaning they remain mostly as the molecule

and only partially ionise.

What we also need to understand:

WATER IS A VERY WEAK ELECTROLYTE

SO THAT MEANS THERE WILL BE HYDROXIDE IONS AND HYDROGEN IONS IN

SOLUTION WHENEVER THERE IS WATER.

That final idea is fundamental to understanding the following methods for determining the

acidity of alkalinity of salts.

Slightly Basic Substances:

If we consider the salt Sodium acetate or NaCH3COO we know:

It will dissociate into ions in water. IN WATER there will also be the hydroxide and hydrogen ions from the water.

So if represent this as two equations one above another:

ZnO(s)+2OH!

(aq)+ H2O(l) "Zn OH( )42!

(aq)

2Cr(s)+2OH

!

(aq)+6H2O(l)"

2Cr OH( )4!

(aq)+ 3H2( g)

Cr OH( )3(s)

+ OH!(aq) " Cr OH( )4!

(aq)

Cr2O3(s)+2OH!

(aq)+ 3H2O(l) "2Cr OH( )4!

(aq)

H2O! H+ +OH!

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Basically what we need to see is: Because NaOH is a strong electrolyte as soon as those ions react with each other they will

dissociate so there is no change. With the CH3COO-and H+however they will react according to the reverse reaction of:

o

THE CH3COOH THEY PRODUCE HOWEVER IS A WEAK ACID.

o So it will not dissociate fully.

o THE CONCENTRATION OF THE MOLECULE IS FAR GREATER THAN

THAT OF THE H+.

! So the concentration of the H+ion will have decreased.

! SOME OF THE HYDROGEN ION HAS BEEN USED UP IN REACTION.

So we get the statement [OH-]>[H+] so pH>7 and therefore the salt is basic.

So we can say that dissolving salts of weak acids in water will form a basic substance but it is better

to understand how to work this out.

Slightly Acidic Substances:

Lets consider this time ammonium chloride NH4Cl, and as before: It will dissociate into ions in water. You will also have the weak ion of water in the solution as well.

Once again looking at the two equations:

So if we are looking at this again we can say: There is no reaction between the Cl-and H+because they form a strong electrolyte. However the ammonium and hydroxide react in the reverse of:

o

ONCE AGAIN THIS PRODUCT THEY PRODUCE IS FAVOURED IN THE

REACTION:

CH3COOH(aq)! H+

(aq) +CH3COO!

(aq)

NH3(g) +H20(l)! NH4+

(aq) +OH!

(aq)

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o So it does not dissociate fully backwards.

o SO SOME OF THE HYDROXIDE IONS HAS BEEN USED UP IN THE

REACTION.

o The concentration of the hydroxide solution will have decreased.

So we get the statement [OH-]


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Which is a weak acid and will only partially ionise back so

therefore there will be a decrease in the concentration of the

H+ions making it basic.

CH3(CH2)16COONa BASIC once again it will react using up H+ion in water to form

weak acid which does no ionise so hydrogen ion concentration decreases and becomes

basic.

K2SO4 BASIC produces weak acid which partially ionises H+ion concentration

decreases and becomes basic.

HBr ACIDIC bromic acid, must be an acid.

NH3 BASIC ammonia, acts as a weak base, should know.

NH4NO3 ACIDIC NH4+would react with OH-in water to produce NH3+ H2O which

only partially ionises back, meaning there is a decrease in concentration of OH-and

solution becomes slightly acidic.

H2O NEUTRAL equal concentrations of OH-and H+.

NH4CH3COO NEUTRAL the NH4would react to produce an acidic solution but the

CH3COO would react to produce a basic solution so they would cancel each other out

giving a neutral solution.

Finally to finish this section off consider the following equation of stearic acid (weak acid found invegetable oils):

Using this show whether Sodium Stearate is slightly acidic or slightly basic.

CO3

2!+ H

+

" HCO3

!

C17H35COOH! H+

(aq) +C17H35COO!

(aq)

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Basically we just show:

So there is a reaction to produce stearic acid, which will only partially ionise back. Therefore there is a decrease in [H+].

So therefore it is slightly BASIC.

Indicators:

Colours of Indicators:

This is simply a table you need to remember and probably put on your calculator:

Introduction:

Basically in an acid base titration and indicator is used so we can see when the substance has been

neutralized before we go any further we need to clear up two very important definitions:

END POINT

This is the point in a titration where a COLOUR CHANGE occurs. i.e. the neutralisation reaction is deemed to be complete.

EQUIVALENCE POINT

H+!" #$ < OH

%!" #$

Litmus Universal

Indicator

Phenolphthalein Methyl

Orange

Bromothymol

Blue

ACID Blue!Red Red Clear Red Orange

BASE Red!Blue Purple/Blue Pink/Purple Orange Blue

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This is the point at which CHEMICALLY EQUIVALENT AMOUNTS (shown via equation)

HAVE BEEN ADDED. So If we were adding HCl to NaOH then we would have the equation:

o

! So the equivalence point occurs when the same number of moles of HCl as

NaOH have been added.

So If we were adding H2SO4to NaOH then we would have the equation:

o

! So the equivalence point occurs when the ratio of NaOH : H2SO4 is 2 : 1. i.e. there are twice as many moles of base as there are acid.

Indicators have a range of pHs they will change in: PHENOLPHTHALEIN 8.3 TO 10.0 METHYL ORANGE 3.1 TO 4.4

NOW we can use the basic ideasto look at three different types of titrations and look at the type of

indicator you would use and WHY!!

Strong Acid vs Strong Base:

If we consider a strong base and strong acid: Strong base NaOH Strong acid HCl

Because these are both strong electrolytes they will not react with the water so the will ONLY

NEUTRALISE EACH OTHER. THEREFORE ANY INDICATOR THAT CHANGES COLOUR IN pH RANGE 4-10 can

be used.

Which can be represented on a graph as:

H+

+OH!

" H2O

H2SO

4 +2OH

!

" SO4

2!+ H

2O

By Josh Walker

Now INDICATORS which are either WEAK ACIDS or WEAK BASES are used in titrations so

that:

THE END POINT IS AS CLOSE AS POSSIBLE TO THE EQUIVALENCE POINT AND A

pH of 7.

o Note that one drop near the end point can change the pH dramatically.


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Now it gets more complicated.

Weak Acid vs Strong Base:

If we consider NaOH being titrated into a solution of the weak acid CH3COOH we get equation:

WHAT IS IMPORTANT: From what we have looked at we KNOW the CH3COO-will react with the water to form a

slightly basic substance:

o As shown by the reaction below.

SO WHAT WE CAN SAY: When the EQUIVALENCE POINT is reached the substances in product will continue to

react producing a basic substance.

o SO THE EQUIVALENCE WILL OCCUR IN A pH>7.

This can be shown on the following diagram:

SO WHAT WE DO:

CH3COOH(aq) +OH!

(aq) "CH3COO!

(aq) +H2O(l)

CH3

COO!(aq)

+H2

O(l)

"CH3

COOH(aq)

+OH!(aq)

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of21 117 In order to account for this basic result we use AN INDICATOR THAT USES UP SOME

BASE. An indicator which BUFFERS the base so that it will still change at equivalence point.

o An indicator that will change colour in pH of equivalence point given. THE INDICATOR WOULD HAVE TO BE A WEAK ACID to use up base.

So we could use phenolphthalein for this situation.

Weak Base vs Strong Acid:

If we consider the weak base NH3aq)being titrated into a solution of the acid HCl we get equation:

WHAT IS IMPORTANT: From what we have looked at we KNOW the NH4+will react with the water to form the

hydronium ion A SLIGHTLY ACIDIC SUBSTANCE:

o As shown by the reaction below.

SO WHAT WE CAN SAY: When the EQUIVALENCE POINT is reached the substances in product will continue to

react producing a acidic substance.

o SO THE EQUIVALENCE WILL OCCUR IN A pH


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Bonding

Atomic Structure:

The Structure of Atoms:

In this course we consider the Bohr Model of he Atom:

What we need to remember about this model is: The electrons not the nuclei determine the chemical properties of the atom. The electrons, which exist in a cloud of negatively charged particles, occupy most of the

volume of the atom. Atoms of each element have a fixed number of electrons. The electron cloud surrounds a very small, dense, positively charged nucleus.

The nucleus consists of positively charged protons and neutrons with no charge. The overall charge on the atom is zero the number of protons and electrons are equal. If there is more than one proton in a nucleus then there must be an equal or greater number of

neutrons to prevent to protons repelling one another. The number of protons determines the type of atom and hence the element.

We also need to remember our symbols relating to this: A = atomic mass number = number of protons and neutrons. Z = atomic number = number of protons, and hence electrons in the atom.

The most important idea of Niels Bohr is that he defined the atom in terms of the electrons not just

lying anywhere around the nucleus but being locked into fixed orbits which have a different

energy level (n):

Although we will consider orbitals in more detail in a few minutes it is worth remembering and

considering:

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of23 117 Each energy level can contain a maximum number of electrons. THE HIGHEST ENERGY LEVEL (shell) IS CALLED THE VALENCE SHELL AND

CONTAINS VALENCE ELECTRONS. During CHEMICAL REACTIONS only the VALENCE ELECTRONS are involved.

It is worthwhile also considering the periodic table of the elements here: The ROW in the periodic table tells you the energy level that the atoms are in.

THE GROUP in the periodic table tells you how many valence electrons that element has:o Except for helium which only has two valence electrons.

It makes sense from there that we can say: Valence electron shells WANT TO CONTAIN 8 electrons (octet rule).

o Except for helium which wants to contain two. This makes sense because all GROUP EIGHT ELEMENTS ARE UNREACTIVE so they

are in a stable state.

It is these ideas which lead to the more complicated situations that we will look at in this course.

You should remember how these valence electrons can be represented as electron dots but we are not

going into it here, for now this is all you need to know on an intro.

Isotopes:

Isotopes are simply Atoms that have the same number of protons but a different number of

neutrons.

From this we can say: Because they have the same number of protons they will be the same element. Because protons must equal electrons in an atom they will therefore have the same

number of electrons and they will thus be chemically identical. Only their physical properties will change.

Isotopes like C-12 and C-14 are often useful in radioactive applications.

Orbitals:

Although the Bohr model is useful it doesnt quite account for the more complicated multi electron

elements.

That is why we look at a quantum mechanics theory: States electrons are fond in regions of space called orbitals.

o An orbital is literally the statistical area in space where you MAY find an electron.

There are a few interesting points we need to know about orbitals: Each orbital with energy level n can hold n2orbitals:

o So energy level 1 can have 1 orbitals.

o Energy level 4 can have 16 orbitals. Each orbital holds TWO electrons. There exists different types of orbitals s orbitals, p orbitals, d orbitals and f orbitals:

o THE S ORBITALS: Contain only one orbital in an energy level.

o THE P ORBITALS: Consist of three dimensions px, pyand pzand hence Contain three orbitals in each energy level they exist in.

o THE D ORBITALS: These contain five orbitals in each energy level they exist in.

o THE F ORBITALS: These contain seven orbitals in each energy level they exist in.

The important thing to understand about these orbitals is that they come in at different times:

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FOR EXAMPLE: In energy level 1 there can only be one orbital:

o Only the 1s orbital. In energy level 2 there is 22= 4 orbitals:

o You have the ONE 2s orbital.

o And the THREE 2p orbitals.

In energy level 3 there is therefore 9 orbitals:o The one 3s orbital.

o The three 3p orbitals.

o And the FIVE 3d orbitals. And finally in energy level 4 there is 16 orbitals:

o The one 4s orbital.

o The three 4p orbitals.

o The five 4d orbitals.

o And the SEVEN 4f orbitals.

It is important to notice that all of the orbitals of the same type and energy level (i.e. all 3ps) are

of equal energy: So when they will up with electrons they each take a single electron before making a pair

as shown below:

Not like this:

This is known as the aufbau (building up) principle.The final thing we need to look at with orbitals is there relative position to one another:

It is not as simple as filling all the ones, then all the 2s and so on. The orbitals can be displayed as below with lowest energy at bottom.

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On your periodic table you can also see the orbitals.

Having looked at all of this information it would we can now see how VARYING VALENCIES in

TRANSITION METALS are formed.

Lets Consider Mn It has two s electrons, which means it has two valence electrons. This implies it will form an ion of 2+.

It can be shown as below.

HOWEVER THE 3d ORBITALS ARE SO CLOSE TO THE 4p ORBITALS THAT AN

ELECTRON CAN JUMP ACROSS TO THE 3d LEVEL

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As a result of this: The substance would still be Mn because the same number of protons. But now there is only one s level electron. Therefore it will form a valence of +1.

That is all we need on the basics of orbitals for now.

Using Orbitals:

We have already briefly discussed this idea but it is important to understand:

THE NUMBER OF VALENCE ELECTRONS IS THE NUMBER OF S AND P ELECTRONS IN

THE HIGHEST ENERGY LEVEL OF THAT ATOM

So when we are looking for number of valence electrons we only need to look at the p and the s in this

highest energy level i.e. for element Rn: 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p6

o Therefore there are 2+6=8 valence electrons.

There is also a shorthand of representing this: Consider P:

o 1s22s22p63s23p3 This has the same INNER SHELL structure as Neon:

o i.e. 1s22s22p6 So for Phosphorus we can actually represent it as:

o [Ne]3s23p3

These two ideas will often come up in tests and exams.

Ionisation Energy:

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Ionisation energy can quite simply be described as the energy required to remove an electron

from an atom to produce and ion. This is the same as the activation energy in a reaction.

This ionisation energy may be in any form of energyheat, light, chemical etc.

So what happens? FIRSTLY ions have a stable octet of 8 valence electrons.

o Therefore atoms that have two valence electrons would want to LOOSE those twoelectrons in that shell to form a 2+ ion and have a stable octet valence shell.

THEREFORE an electron is removed via the IONISATION ENERGY in the reaction. ANOTHER ELECTRON MAY BE REMOVED but this would take more energy

because the second electron is being removed from a positive ion.

o i.e. If we look at Mg the first electron is removed.

o Then the second you are removing a negative electron from a positive ion so you

have to break the attraction due to opposite charges as well.

o So it takes more energy for the second.

From this it makes sense REMOVING ELECTRONS FROM NOBEL GASES (atoms with

stable octet) TAKES VERY LARGE AMOUNTS OF ENERGY: Because the atoms are already an ideal energy state. So an electron does not want to be removed.

From this simply idea we can use tables and graphs of ionisation energies to predict the valence or

number of valence electrons that an atom has: Basically all you do is look for a jump:

o The ionisation energies will increase in a relatively linear pattern.

o When the atom reaches a stable octet however there will be a big jump in the energyrequired.

o So how many electrons removed before this jump will be the number of valence

electrons in the atom.

So if we consider the following graph:

It is clear that the jump occurs after three electrons have been removed and thus the substance will have

a 3+ valence.

Now lets move on.

Electronegativity Introduction:

Mg+IE!Mg+ + e"

Mg+ +IE!Mg2+ + e"

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Electronegativity can basically be defined as the ability of an atom to attract electrons.

Why does this occur?

o Basically atoms in the higher groups have 6 or 7 valence electrons.

o THEY THEREFORE WANT TO GAIN 1 OR 2 ELECTRONS to form a stable octet and

hence ion.

o So they are attracting electrons.

We will look at trends of electronegativity in more detail later but for now it is worth knowing.A order of ionisation energies for this highest can be:

Now lets jam this together with ionisation energy to make something worthwhile.

Formation of Ions:

We have already looked at how metals (substance in lower groups of periodic table) LOOSE

electrons to form positive ionswhilst non-metals (substances in higher groups of periodic table),

GAIN electrons to form negative ions.

We can happily say:

IONS HAVE THE SAME VALENCE ELECTRON STRUCTURE AS A NOBLE GAS 8 valence

electrons except for those near helium that have 2 valence electrons.

But atoms do not really LOOSE electrons what happens can be best described by an example

consider Sodium and chlorine:

o Sodium has the following electron dot structure.

o Chlorine has the following dot structure.

o If the two were to come near each other:

o ENERGY with energy from surrounding eventually there would be enough remove

an electronfrom the sodium.

o CHLORINE - would attract this electron and pick it up also forming a stable octet.

o SO the atoms form ions via EXCHANGING electrons.

We can also describe ionic bonds using this logic:

o The two atoms collide.

o AN ELECTRON IS TRANSFERRED FROM THE SODIUM TO THE CHLORINE.

o Both atoms now form ions.o These ions are magnetically attracted to each other as shown below:

F>O> Cl > N> S> C> H

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o The arrow shows electron going from the sodium to the chlorine.o You do not draw the electron dots on the sodium simply because it makes it look

complicated.

The only other thing you need to remember about ions is:

o POSITIVE IONS are often called cations.

o NEGATIVE IONS are often called anions.

That is the end of basically the revision with a few extra ideas added in. Now we can move into some

partially new information.

Chemical Bonding and Properties of Materials:

We know that the chemical properties of a substance are dependent on the number of valence electrons.

However the physical properties are dependent on the type of bond between the atoms.

o Bonds may be:

o Ionic.

o Metallic.

o Covalent.

! Covalent molecular.

! Covalent Network.

o The SHAPE and POLARITY will determine the type of bonds between individual

(DISCRETE) molecules.

o Melting points and boiling points are an indication of the strength of the chemical bond

(force).

It is also important to remember that the physical properties we are looking at are:

o Melting points.

o Boiling points.

o Hardness.

o Conductivity of electricity and heat.

o Solubility.

So lets move in and look at each of the type of bonds above separately.

Ionic Bonding:

When the metals form positive ions, and the non-metals negative ions these +ve and ve ions attract

electrostatically forming a 3-dimensional crystal (lattice, network) structure.

THIS ATTRACTION IS CALLED AN IONIC BOND

Ionic solids are :

o Hard.

o Because of network structure.

o Brittle.

o Because of positive and negative ions if you bend it then they will align and repel.

o

Non-conductors of electricity.o Because there are no free ions or electrons to conduct a current.

o High Melting Points.

o Because there is a strong bond in the crystal lattice.

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When they are dissolved in water or melted however they do conduct electricity:

This is because they break up into ions!!

o Consider NaCl.

o When melted or dissolved it dissociates into ions as shown below:

o The now mobileions are attracted to oppositely charged electrodes and hence conduct

a current.

You can also have complex ions which remain together when dissolvedthese include:

o NO3-

o NH4+

o They have a covalent bond between the atoms within the ion.

So substances like NH4Cl have both ionic and covalent bonds within them.

Metallic Bonds:

Metallic bonds are different to any other sort of bonds.

They form a crystal structure (lattice, network) with an array of positive metal ions surrounded

by a sea of mobile delocalised electrons not belonging to any particular atom.

This can be seen below:

As a result of this metals have:

o Varying melting points.

o Due to varying number of protons etc in bonds.

o Vary in hardness.

o Due to varying number of protons etc in bonds.

o Malleable and Ductile.

o Because delocalised electrons ensure there is always attraction and bending does notalign them and repel.

o Shiny when polished have lustre.

o Conduct heat and electricity well.

o Because mobile means moving so the electrons can move freely and conduct a current

and/or heat.

Metals are really quite simple to understand in terms of bonding.

Transition Metals:

These group of metals are Metals in which the d orbitals are filled with electrons.

o They usually have 2 valence electrons the two s electrons.o They have all the typical metal characteristics.

o They also have some special characteristics which include:

NaCl! Na+

+Cl"

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of31 117o Variable valencies.

! Eg. Fe2+and Fe3+all exist.

o They Form Coloured Compounds.

! Cr2+- blue.

! Cr3+- green.

! CrO42-- yellow.

!

Cr2O72-- orange! MnO4-- purple

! Cu2+- blue

o For Complex Ions.

Complex ions form when a positive transition metal ion is SURROUNDED by anions (Cl-, CN-)

or polar molecules (H2O, NH3).

You can say these ligands are bonded to the central ion by coordinate covalent bonds:

We need to know a few reactions for the course so lets just summarise them here:

Ammonia solution is added to copper II sulfate solution:

Ammonia solution can be seen as a mixture of NH3, H2O, NH4+and OH-because:

Therefore it make sense the reaction would be:

Excess ammonia solution is now added to this copper precipitate:

Excess ammonia is added to the initial copper II sulfate solution:

Zinc reacting in the same fashion as above:

Silver also reacting as above:

Now we know these reactions we can deduce other onesi.e.A solution of Cobalt Nitrate reacts with

water to form the complex ion hexaaquocobalt II [Co(H2O)6]2+.

o Since it only reacts with water it follows that with a bit of logical balancing we can get:

Thats the extent of transition metals and complex ions in this course.

NH3+ H

2O! NH

4

+

+OH!

Cu2+(aq)+ 2OH!

(aq) " Cu OH( )(s) Pale Blue ppt

Cu OH( )2( s)

+ 4NH3(aq)! Zn NH3( )4"# $%2+

(aq)+2OH&(aq)

Cu2+(aq)+ 4NH3(aq)! Cu NH3( )4"# $%2+

(aq)

Zn2+(aq)+ 4NH3(aq)! Zn NH3( )4"# $%2+

(aq)

Ag+(aq)+2NH3(aq)! Ag NH3( )2"# $%+

(aq)

Co

2+

(aq)+6H2O(l)! Co H2O( )6"# $%2+

(aq)

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Covalent Molecular Bonding:

Covalent bonds form when non-metal atoms share electrons and produce a discrete (individual) unit

called a molecule.

COVALENT BONDS are bonds formed by SHARING ELECTRONS and they are the

STRONGEST TYPE OF CHEMICAL BOND between ATOMS

However the bonds BETWEEN these individual units called MOLECULES are WEAK.

o The type of the bond between the molecules is determined by how the molecules arearranged and determines the strength and type of substance you can expect.

Some properties of covalent molecular bonded substances are:

o Brittle OR soft when solid.

o Solids have low melting points.

o Liquids have low boiling points.

These last two points bring out a very important idea for the properties of these materials:

o It suggests that there is a weak force between the molecules of the substance.

o However there is still a strong force between the atoms in the molecules.

So:MELTED BOILED OR DISSOLVED they remain as molecules

For example:

When sugar dissolves it breaks away from the crystalbut remains as individual molecules in

the water.

The same happens for water as it boils they overcome the molecular bonds between the

molecules but the atoms within the molecules stay bonded.

All this brings up the final property of covalent molecular substances:o THEY DO NOT CONDUCT ELECTRICITY OR HEAT.

o This is because all the electrons are tied up in the covalent bonds and there are no

ions so therefore there is nothing to carry the current or heat.

That is the basics of covalent network solids for now.

Covalent Network Solids:

There are a few substances that form three dimensional crystalline networks of covalently bonded

atoms rather than indicvisual molecules.

o Carbon (diamond) C

o Silicon Si.o Silicon Carbide SiC.

o Silicon Dioxide (quartz) SiO2

C12H22O11(s)H2O

! "!! C12H22O11(s)

H2O(s)! H2O(l)! H2O(g)

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of33 117o Boron B

Carbon also forms a two dimensional network:

o Graphite C.

Because of its two dimensional nature graphite has some characteristics different to that of the other

covalent network solids:o Soft.

o Because there are weak bonds between the carbon layers.

o High Melting Point.

o Because there are strong covalent bonds between the carbon atoms within the layer.

o Conducts Electricity.

o Because each carbon is only surrounded by THREEother carbons there is therefore a

free delocalised electronfor each carbon.

o This electron can move through the layersand conduct electricity.

The properties for the other three dimensional covalent networks solids are all similar:

o High Melting Point.o Hard.

o Non Conductor.

o This is all because of the strong covalent bonds that holds all electrons together,

making it hard to break apart, hard to touch and a non conductor.

Now we have completed the look at bonding properties we can move into the section of total new

work.

Shapes and Polarities of Molecules:

We have for a while now looked at molecules and covalent bonding and discussed the idea of polarity

but now it is time to officially understand these terms and use them.Before we go any further we need to understand A molecule is said to be POLAR when it has a

dipole that is a positive and a negative end.

We will look at how that occurs and the different ways it can occur in this section.

Linear Non-Polar:

If we consider the chlorine molecule Cl2.

o Each chlorine atom has 7 valence electrons so it wants to share one with the other chlorine

molecule.

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o As you can see there are both bonding and non bonding electron pairs.

This is a LINEAR NON POLAR molecule:

o The Molecule is shaped in a straight line LINEAR.

o The electrons are equally attracted to both nuclei.

o Therefore they are evenly symmetrically distributed so there is NO DIPOLE.

This is the same with other halogens F2Br2I2and also H2

Linear Polar:

If we consider the hydrogen chloride molecule HCl.

o The hydrogen atom has one valence electron whilst the chlorine has seven so they both

want to gain one therefore they share one.

o As you can see there are both bonding and non bonding electron pairs.

This is a LINEAR POLAR molecule:

o The Molecule is shaped in a straight line LINEAR.

o Chlorine is more ELECTRONEGATIVE so inside the nuclei the electrons are attracted to

it more than they are to the hydrogen.

o There are also a far greater number of non bonding electron pairs on the chlorine atom.

o Therefore the molecule has a slightly negatively charged end and hence a slightly

positively charged end.

We represent this with "- (for negative end) and "+ (for positive end) with a line pointing in the

direction of the negative end with a cross over the positive end

This is the simplest type of polar molecule.

Bent, Polar:

If we consider water H2O

o There are two hydrogen atoms attached to the one oxygen atom.

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o Oxygen is more electronegative than Hydrogen.

o There are also free non-bonding electrons on the oxygen.

o So we have two dipoles as shown below.

o This adds up to give one dipole.

This is a BENT POLAR molecule:

o The Molecule is BENT because the bonding and non bonding electron pairs REPEL so

there is an angle between them.

o DIPOLE EXISTS because of oxygen being more electronegative than hydrogen and of

excess non bonding electron pairs.

o The two dipoles add up to produce a NET DIPOLE.

Other molecules like this include H2S, F2O, Cl2O

Tetrahedral Non Polar.

Considering methane, CH4:o To The one carbon atom there are four hydrogen atoms.

o Carbon is more electronegative than hydrogen.

o However the hydrogens are all in opposite directions so you will have four dipoles.

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o These dipoles all cancel so the electrons are evenly distributed and we have a non polar

molecule.

This is a TETRAHEDRAL NON POLAR molecule:

o TETRAHEDRAL because bonding electron pairs and H nuclei repel so there is a

maximum distance between the C-H bonds (angle 109.5)spread out.

o NON POLAR because the four dipoles all cancel each other out an end to give a overall

non polar effect..

Other molecules like this include CCl4, ALL straight chain HYDROCARBONS, SiH4(silane)

Tetrahedral Polar:

Considering CH3Cl:

o To The one carbon atom there are three hydrogen atoms and one chlorine atom.

o Carbon is more electronegative than hydrogen and chlorine is more electronegative than

carbon.

o So we have three dipoles (caused by the hydrogen), which are the same, and one dipole

(FROM THE CHLORINE) WHICH IS DIFFERENT.

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o SO THE DIPOLES DO NOT CANCEL.

This is a TETRAHEDRAL POLAR molecule:

o TETRAHEDRAL because again there is repulsion between the bonding electron pairs

and nuclei so they spread out.

o POLAR because the four dipoles do not cancel, because they are not all of the same

strength

Most molecules in form XY3Z will be polar tetrahedral molecules

It is worth considering the case CH2Cl2:

o THE IMPORTANT POINT IN THIS SET UP IS REMEMBERING THAT UNBOUNDELECTRON PAIRS ARE ALWAYS LEFT TOGETHER:

o So when the hydrogen fills up two of the carbon electrons.

o The remaining two places are NEXT TO EACH OTHER.

o As a result the structure is like this:

NEVER LIKE THIS:

o So from the correct structure it is easy to see how there is a dipole.

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This is a TETRAHEDRAL POLAR molecule:

o TETRAHEDRAL because again there is repulsion between the bonding electron pairs

and nuclei so they spread out.

o POLAR because the four dipoles do not cancel, because they are not all of the same

strength

It is incredibly important you remember this structure note

Triangular Pyramid, Polar:

Considering ammonia, NH3:

o To The one nitrogen atom there are three hydrogen atoms.

o Nitrogen is far more electromagnetic than hydrogen.

o

The Nitrogens are in three directions so you have three separate dipoles.

o Because of the repulsion between the electrons the Hydrogens spread out to form a

triangular pyramid shape and this REINFORCES THE DIPOLE.

This is a TRIANGULAR PYRAMID EXTREMELY POLAR molecule:

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of39 117o TRIANGULAR PYRAMID because bonding electron pairs and H nuclei repel so there

is a maximum distance between the C-H bonds and this forms a triangular pyramid.

o POLAR because the three separate dipoles all reinforce one another to make it a very

polar molecule.

Other molecules like this include most in the form of XY3

Special Substances:There are a few cases we need to consider some questions some just molecules:

Carbon Dioxide:

In carbon dioxide:

o A double bond exists between the carbon and oxygen:

o The Oxygen is more electronegative than the carbon so there are two dipoles which exist.

o These dipoles cancel however resulting in a LINEAR NON POLAR MOLECULE.

So although from looking at this CO2formula we may think it is the same as H2O if we actually THINK

about what is happening with the bonds we will see that it is definitely not the case.

A Polar molecule that has a non polar bond within the molecule:

This is a common question in tests and exams and requires a simple answer hydrocarbon with more

than one carbon atom:

A Non Polar molecule that has a polar bond:

This is identical to the question above there are polar bonds in all the hydrogens attached to a carbon

atom so you just need a hydrocarbon:

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If you are able to understand and do this then you have a good chance of success:

Electronegativity and Polarity Strength:

We have already briefly looked at this but it is worth clarifying.

The greater the difference in electronegativity in a molecule the more polar the molecule will be:

So therefore:o Cl-Cl no polar, because of zero difference in electronegativity.

o C-Cl very polar because of large difference in electronegativity.

o IONIC BONDS extremely polar because electrons are actually transferred so you have

large difference in electronegativity.

This is not the most important concept but it is useful if you need to identify the most polar molecule

in a series of molecules.

Co-ordinate (dative) Covalent Bonding and Ions:

A co-ordinate (dative) bond is simply a covalent bond where one atom donates bothe electrons to

the bond.For example in the ammonium NH4+ion:

Ammonia reacts with a proton to form the ion:

AS YOU CAN SEE the proton latches on to the two free electrons of the ammonia and thus the

ammonia has provided both electrons for the bond and it is a dative bond.

Dative bonding is the basis for many of the complex ions such as:o H3O+

o SO42-

o NO3-

o And so on.

It is very important to remember:

ALL IONS ARE NON POLAR!!!!

This is because in order to be polar a molecule must have a dipoleand ions only have EITHERa

positive OR a negative end.

It is essential to understand all of this before we move on to the next section.

Forces Between Molecules Intermolecular Forces:

It is important that we understand there are bonds withinand bonds betweenmolecules:

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of41 117o Covalent bonds exist WITHIN molecules (between the atoms)

o This is the strongest chemical bond that exists.

o Van der Walls forces exists BETWEEN molecules:

o These molecular forces are very weak.

This weakness can be easily seen if we compare the melting points of various solids (REMEMBER

MELTING POINT IS AN INDICATION OF BOND STRENGTH):

As you can see the bond strength varies even in molecules. This depends on the polarity and shape of

the molecules. ALSO NOTE: There are virtually NO intermolecular forces between particles in a

GASEOUS state.

We can display van der walls forces in a chart below:

Now lets investigate each of them.

Dispersion Forces:

Dispersion forces ARE PRESENT BETWEEN ALL MOLECULES!!!!!!!!

The important point however is that:

o THEY ARE THE ONLY INTERMOLECULAR FORCES BETWEEN NON POLAR

MOLECULESin a liquid or solid state.

The logic behind the dispersion forces are very simple:

o Electrons can always move around inside atoms.

o OCCASIONALLY THESE ELECTRONS WILL ALIGN TOGETHER IN ONE

DIRECTION.

o THIS CAUSES ONE MOLECULE TO HAVE A TEMPORARY DIPOLE.

o Because it is next to another molecule this will induce a TEMPORARY DIPOLE in that

molecule as well.

o And as a result there is a force of attraction between the molecules.

It does not take a genius to work out that a force of attraction dependent on the random motion of

electrons is A VERY WEAK FORCE the weakest of all intermolecular forces.

METALLIC IONIC COVALENT

NETWORK

MOLECULAR

(POLAR)

MOLECULAR

(NON POLAR)

1100C 800C 1700C 0C -80C

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It is also important to note AS MOLECULAR WEIGHT INCREASES SO DOES MOLAR MASS:

This occurs because:

o An increase in molecular weight means there are more protons and electrons in the atom.

o

Therefore there is a greater force of attraction in the molecules.o Thus you need more energy to break the molecules apart.

Substances that are bonded by dispersion forces are:

o Noble gasses.

o Halogens.

o Hydrocarbons (alkanes, alkenes,

alkynes).

o Waxes.

o CO2.

o CS2.o SiH4.

o CCl4.

o S8.

o P4.

o O2.

o N2.

o H2.

Thats as simple as dispersion forces are.

Dipole-Dipole Forces:

Dipole-Dipole forces EXIST BETWEEN POLAR MOLECULES in a liquid or solid state!!!

The logic here is the simplest:

o Basically every molecule has positive and negative end.

o So there is an attraction between the +ve end of one molecule and the ve end of another.

This can be shown below using HCl in liquid form as an example:

Because there is a set electromagnetic attraction dipole-dipole forces are STRONGER than

dispersion forces.

Substance which are dipole-dipole bonded include:

o SO2.

o NO2.

o HBr.

o HI.

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of43 117o H2S.

o Alkyl Halides (CH3Cl, CH3CH2Br

etc).

These are the easiest intermolecular forces to understand.

Hydrogen Bonding:

A PARTICULARLY STRONGdipole-dipole force that exists in some liquids and solids is AHYDROGEN BOND!!

o Its strength is due to the fact that Hydrogen does not have many electron clouds, the electrons

are close to the nucleus and hence there is a stronger force of attraction there.

We can say:

Hydrogen Bonds Occurs in molecular substances that contain POLAR MOLECULES with

HYDROGEN BONDED TO:

o OXYGEN.

o NITROGEN.

o FLUORINE.

H - F O NCommon substances include:

o H2O.

o NH3.

o HF.

o CH3OH.

o CH3CH2OH.

o CH3COOH.

You can basically say that THE PRESENCE OF THE OH GROUP INDICATES HYDROGEN

BONDING.

If we were to draw out a hydrogen bond it would look pretty much exactly the same as a dipole-dipolebond:

Hydrogen bonding can be used to explain the following two graphs:

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Simply, the molecules of H2O and HF are held together by hydrogen bonds. The molecules of

the other substances are only held together by dipole-dipole bonds. The hydrogen bonds are

much stronger and so therefore although the molecular mass is lower (boiling point increases

with molecular mass) the boiling point is higher.

And thats all we have to say about hydrogen bonding.

Explanations:The most common question in an exam is explain why so here we will look at a few examples

of this explain scenario:

EXAMPLE ONE

Water (M=18g) boils at 100C and Methane (M=16g) boils at -161C. Although the molar masses are

similar there is a large difference in boiling points explain:

Water is a polar molecule:

The intermolecular forces in water are a particularly strong dipole-dipole bond called a

hydrogen bond. This is the bond between the oxygen and the hydrogen.

Methane CH4on the other hand is a non-polar molecule. Therefore the molecules of

methane are only held together by weak dispersion forces that is forces of attraction

caused by the alignment of electrons in molecules. As a result of this the energy required

to separate the methane molecules is far less than that required to separate the water

molecules.

EXAMPLE TWO

The boiling points of the first 4 alcohols are:

o Methanol, CH3OH - 64C

o Ethanol, CH3CH2OH - 78C

o Propanol, CH3CH2CH2OH - 98C

o Butanol, CH3CH2CH2CH2OH - 118C

Explain why this occurs:

Simply, the boiling point is increasing because the molar mass is increasing. This means

there are more protons and more electrons in each atom. Therefore the force of attraction

between the atoms in the bonds is stronger and it takes MORE ENERGY and hence a

greater temperature to break them apart.Thats everything for the intermolecular forces in this course.

Solubility:

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Solubility depends on the interactions between the bonds in the solute and the solvent:

o If the force between thesolvent-soluteis stronger than the force between thesolute-solute or

solvent-solvent then a substance will dissolve.

So lets look at a few ways things dissolve and understand it all.

Dissolving Ionic Solids:

The process behind the dissolving of ionic solids is possibly the simplest and most logical to explain which is why we begin here.

Simply we can say:

o When an ionic solid is placed in water it has a negative end and a positive end.

o The water molecules are also polar.

o Therefore:

o The positive end of the water molecule will attach to the negative ion.

o The negative end of the water molecule will attach to the positive ion.

o The subsequent KINETIC ENERGY (MOTION)of the water molecules gradually

BREAKS UP the crystal.

o Thus the ionic solid is now separated into IONS and each ion is surrounded by a watermolecule making it soluble.

It is important to consider:

AN IONIC SOLID IS INSOLUBLE WHEN THE ATTRACTION BETWEEN THE IONIC

BONDS IS GREATER THAN THAT BETWEEN THE IONS AND THE WATER.

It is also worth noting:

o SOLUBILITY OF SOLIDS IN LIQUIDS increases as temperature increases.

o Solubility of gasses decreases in liquids as temperature increases.

Now lets look at molecular substances and see how they dissolve.

Dissolving Polar Molecular Substances:

There are two points we need to consider about polar molecular substances:

o THEY DISSOLVE IN OTHER POLAR SUBSTANCES USUALLY:

o Because there is an attraction between the negative end of one molecule and the

positive end of another molecule.

! This is particularly strong when the solute forms hydrogen bonds with the

solvent.

! i.e. the following dissolve by forming hydrogen bonds with water.

Methanol (and other alcohols). Ammonia.

Glucose.

! Remember that although they have dissolved molecules are still molecules they

do NOT dissociate like ionic solids.

o HOWEVER, if the bonds between the solute are stronger polar bonds than that

which would form between the solute and the solvent (i.e. hydrogen bonds) then

they would not dissolve.

o THEY DO NOT DISSOLVE IN NON POLAR SUBSTANCES:

o This is simply and logically because the dipole-dipole OR hydrogen bonds between

the solute are far stronger than the dispersion forces in the solvent.

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of46 117! So there is NO BOND which could possibly be strong enough to break the

polar substance up.

We will look at this more later but long chain hydrocarbons attached to an OH at the end i.e.

CH3CH2CH2CH2CH2CH2CH2CH2CH2CH2CH2OHusually are only slightly soluble because the non

polar hydrocarbon end MASKS the polar OH end.

And thats all there Is to the dissolving of polar molecular substances.

Dissolving Non Polar Molecular Substances:

Once again with non-Polar molecular substances we need to again consider two scenarios:

o THEY DISSOLVE IN NON POLAR MOLECULAR SUBSTANCES:

o This is simply because there is dispersion forces in the non polar solvent.

o When the molecules of the solute enter into it they will be affected by the dispersion

forces.

o Therefore temporary dipoles will be induced in them as well and they will dissolve.

! Common examples of this include:

Oil C16H34dissolving in petrol C8H18.

Iodine I2dissolving in tetrachloromethane CCl4.

o THEY DO NOT DISSOLVE IN POLAR MOLECULAR SUBSTANCES:

o This is simply because they cannot form any electromagnetic bonds with the water.

o SO THE WATER-WATER or (POLAR-POLAR) hydrogen or dipole-dipole bonds

are far stronger than POLAR-NON POLAR bond.

o The polar bonds overpower the dispersion forces of attraction between them.

Thats all there is to non polar substances.

The Rule:

So these two investigations above will lead us to a useful rule (which although not good for explainingcan help for identifying):

LIKE DISSOLVES LIKE

Polar dissolves polar

Non polar dissolves non polar

Solubility is really that simple.

Ionisation and Hydrolysis:

Sometimes polar molecules form such strong bonds with the water that the covalent bonds is

broken and the molecule IONISES:

e.g. the formation of the hydronium ion.

Another important idea is when polar molecules dissolve by REACTING (HYDROLYSIS) with

water forming acids and bases:

e.g. reaction of ammonia.

Thats all there is on solubility in this course for now.

Periodic Table Summary:

The periodic table gives many trends so it is worth summarising them here and understanding themso they can be used in an examination.

HCl+ H2O! H

3O

+

+Cl"

NH3+ H

2O! NH

4

+

+OH"

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Ionisation Energy:

In two parts:

INCREASES ACROSS A ROW IN THE PERIODIC TABLE

Because:

o The number of valence electrons increases.

o Therefore there is a greater attraction between the nucleus and the electrons.

o So it takes more force and more energy to remove an electron.

DECREASES DOWN A GROUP IN THE PERIODIC TABLE

Because:

o There is an extra energy level of electrons each time.

o Therefore the space between the protons and the valence electrons is greater.

o Thus the force between them decreases.

o This makes it easier to remove an electron from the valence shell.

Electronegativity:

In two parts:


Because:

o The number of electrons is getting closer and closer to a stable octet,

o Therefore it just wants to gain electron/s to reach that stable octet.

o So it attracts electrons with much more strength.

o NOTE: THE NOBEL GASSES HAVE A IONISATION ENERGY OF ZERO BECAUSE

THEY HAVE A STABLE OCTET AND THUS WILL NOT ATTRACT ELECTRONS AT

ALL.

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DECREASES DOWN A GROUP IN THE PERIODIC TABLE

Because:

o There is more space between the protons and the valence electrons.

o Thus there is less force of attraction between the valence electrons and the nucleus.

o Thus the atoms are less likely to attract an electron.

Atomic Size:

In two parts:

DECREASES ACROSS A ROW IN THE PERIODIC TABLE

Because:

o Because the number of valence electrons increases.

o Thus the force of attraction between the nucleus and valence electrons also increases.

o This brings the valence electrons closer to the nucleus thus decreasing the size of the atom.

INCREASES DOWN A GROUP IN THE PERIODIC TABLE

Because:

o There is always an extra energy level (series of orbitals) between the valence electrons and the

nucleus.o Thus the atoms must obviously increase in size simply because there is another layer of stuff.

Oxidising/Reducing Agents:

Before we look at this we need to revise oxidation and reduction:

o OXIDATION is where the substance LOOSES ELECTRONS.

o REDUCTION is where the substance GAINS ELECTRONS.

o A REDOX reduction-oxidation REACTION is where the loss and gain of electrons

attract.

o ALL REACTIONS OF ELEMENTS ARE REDOX.

Knowing this it makes sense to sayThe ease of reductionor strength as an oxidising agent:


Because:

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of49 117o Because electronegativity also increases.

o So it becomes easier for the substance to gain electrons.

o And thus it becomes easier for it to be reduced and substances that are reduced are good

oxidising agents.

The ease of oxidationor strength as a reducing agent:

INCREASES DOWN A GROUP IN THE PERIODIC TABLE

Because:

o Because the ionisation energy is also decreasing.

o So it becomes easier for the substance to loose electrons.o Thus it becomes easier for it to be oxidised and oxidised substances are reducing agents.

That is as far as we need to go in bonding for now.

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Electro Chemistry

Electrochemical Cells:

We can safely say that:

Electrochemical cells convert chemical energy into electrical energy

They do this via a chemical REDOX reaction

But as with most things,simple definition will not get you far we need to have a much deeper

understanding of what is happening, so what do we know:

o Redox reactions involve the transfer of electrons.

o An electrical current is the flow of electrons.

o So there seems to be a connection there.

Also we need to remember that redox reactions occur due to the varying ability of substances togain or loose electrons:

o The abilities are shown on the E0table:

o Given a reaction on the left side anything below it will react.

o Given a reaction on the right side anything above it will react.

To understand the logic behind electrochemical cells we need to consider a piece of zinc metal being

placed in a solution of copper sulfate:

o The zinc metal will be oxidised to produce Zn2+ions.

o The Cu2+will be reduced to form copper metal.

o This can be shown on the diagram below:

o The reaction of this will be:

The idea for an electrochemical cell is now very simple to understand:

o If you could spit this reaction up so the electrons transferred had to travel along a wire

to complete the reaction then you would be generating a current.

Reduction:Cu2+(aq) + 2e!

"Cu(s)

Oxidation:Zn(s) "Zn2+

(aq) + 2e!

Zn(s) +Cu2+

(aq)! Zn2+

(aq) +Cu(s)

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of51 117! Because a current is a flow of electrons:

o And voila, below is the first electrochemical cell:

! The reactions have been separated so electrons travel along a wire (hence we

have a current)

! The circuit is completed with a salt bridge containing some ionic substance.

It is important however that we look at how exactly this works and understand the logic behind itfurther.

Introductory Logic:

To understand lets consider a more proficient diagram of the electrochemical cell above:

Above we have:

o Two half cells:

o One where reduction occurs.o One where oxidation occurs.

! The Cu metal electrode in Cu2+electrolyte is the Cu/Cu2+half-cell.

! The Zn metal electrode in the Zn2+electrolyte is the Zn/Zn2+half-cell.

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of54 117! They therefore move through the solutions ensuring the current is carried on

in the form of the ions.

The last point of importance is THE EXPECTED VOLTAGE:

o We know from our previous topic of REDOX that a redox reaction has a voltage potential.

o Therefore the EXPECTED VOLTAGE of any cell will be the two E0values for each

half cell added together.Looking at essential basics this is about all we need to cover for now.

Step-by-Step Worked Example:

Now we have got the basics it is worthwhile considering an example and working through it step by

step so lets consider the following question:

A half cell with a strip of manganese metal is placed in manganese chloride solution and is

joined to a cobalt metal strip placed in cobalt nitrate solution via a salt bridge:

Determine:

*Anode and cathode.

*Positive and negative electrode.*Sites of Oxidation and Reduction.

*Direction of electrons moving through external circuit.

*Direction of ions moving through salt bridge.

*Half equations at each electrode, and overall redox equation with expected voltage.

STEP ONE: Get the Diagram

The first thing we need to do is make sure we have a clear diagram of what we are looking at in

front of us:

STEP TWO: Identify What You Have

The second step is to consider what we have which is in this case:

*Mn.

*Mn2+.*Cl-.

*Co.

*Co2+.

*NO3-.

Now you locate these on the diagram:

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STEP THREE: Use E0table to work out what is oxidised and what is reduced and hence which

electrode is your anode and which is your cathode:

Now you basically look for whether the Co or Mn will be more easily oxidised and which ions

will be more easily reduced and you work out what reaction will occur at each elecrode.

From there you can use standard definitions to solve:

STEP FOUR: Now Find Movement of Ions, Direction of Electron Flow, +ve and ve etc:

Use the rule that cations go to the cell with the cathode and anions go to the cell with the anode

to determine the movement of the ions.

Electrons flow from negative to positive.

By definition we know the anode is negative and the cathode positive.

Draw this all on the diagram:

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STEP FIVE: Combine to one Redox equation and find predicted voltage E0:

We can say:

Following these steps should enable you to answer any electrochemical cell question.

Inert Electrodes:

You do not always need reactive metal electrodes to produce an electrochemical cell:

Platinum Electrodes are Chemically Inert However:

! The Hydrogen Will be oxidised.

! The Chlorine will be reduced.

! They will produce a cell with expected voltage 1.36V.

Reduction: Co2+(aq) + 2e!

"Co(s) E0

= !0.28V

Oxidation:Mn(s) "Mn2+

(aq) + 2e! E0 = +1.18V

Redox:Mn(s) +Co2+

(aq)!Co(s) +Mn2+

(aq) E0= 0.9V

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NOTE: Because there are no solutions, which would mix, these share a common electrolyte of

HCl.

And that is pretty much all there is to electrochemical cells in the straightforward manner.

Commercial Cells:

Although we may know how cells operate it is important to understand how they are produced and

used in the commercial environment. Below therefore are examples of the three most common types of

electrochemical cells and an explanation about how they work.

Dry Cells:

This is the most common cell and is used inn torches, as well as other electrical appliances. The basic

structure can be seen on the diagram below:

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As you can see there are a few essential elements to the structure:

o OUTER ZINC COATING:

o By looking at E0table we can see this is the anode (-).

o Therefore the zinc will oxidised according to:

!

o GRAPHITE ROD SURROUNDED BY MANGANESE DIOXIDE PASTE:

o This rod is the Cathode (+)o The manganese around it in the paste is reduced making the cathode reaction as

follows:

!

o ALUMINIUM AND ZINC CHLORIDE IN PASTE:

o These are both ionic substances, which will break into ions in solution.

! Hence they are the Electrolyte for the reaction.

Also the ammonium ions provide the hydrogen ions needed for the

cathode reaction via:

From the reactions above we can now identify that:

o The cell has potential maximum e.m.f of 1.48V.

Also another VERY IMPORTANT POINT THAT:

o The cell is a primary cell that is once it has been discharged (reaction complete) it cannot

be recharged and therefore must be discarded.

Another interesting idea surrounding the dry cell is the ALKALINE VERSION of the dry cell that

has a superior performance to the acid form:

o A powdered Zinc anode is used instead.

o The electrolyte is KOH.

Zn(s) !Zn2+

(aq) +2e"

2MnO2( s) +2H+

(aq) +2e!

"Mn2O3(s)H2O(l)

NH4+

(aq)!

NH3(aq)+

H+

(aq)

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of59 117o Basically this type of cell will work for longer and supply current more rapidly

however it is also more expensive.

So if we were to look at the reactions taking place:

o ANODE:

o

o CATHODE:

o

or (depending on what substances are used in the cell)

o

Thats all there is for the simple dry cell.

Wet Cells:

Also known as the LEAD ACID BATTERYthis cell is commonly used in motor vehicles. The

standard characteristics include:o A number of cells each cell has a voltage of approx 2.0V.

o Inside each cell there is:

o Heavy-duty external casing (on outside).

o Bank of lead grids supporting a large surface area of the electrode material:

! Negative electrode grid is filled with spongy metallic lead.

! Positive electrode grid is filled with brown lead dioxide.

o Electrolyte used is sulfuric acid.

This can be seen below:

The main difference in this cell is that:

o THE CELLS IN THIS BATTERY CAN BE RECHARGED therefore it is described as a

secondary cell.

If we consider the reactions occurring we can say:

Zn(s)+2OH!

(aq) "Zn OH( )2( s)+2e!

Ag2O(s) +H2O(l) +2e!

"2Ag(s) +2OH!

(aq)

HgO(s) +H2O(l) +2e!

"Hg(l) +2OH!

(aq)

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Now what we need to consider:

o As the batter is discharged, electrodes become coated with insoluble lead sulfateand the

sulfuric acid is used up:

o Thus the density of the electrolyte decreases.

! Hence the degree of charge of battery can be checked by testing electrolyte

density.

Basically this battery can be recharged simply by connecting the terminals to another electrical

source of higher voltage and reversing the current through the circuit.

o Hence the reaction is reversed and we have:

!

Once again, in this battery electrical energy is STOREDas chemical energy in the RECHARGING

PROCESSbefore it is AGAIN RELEASED AS ELECTRICAL ENERGY.

And that, is all there is to wet cells.

Ni-Cad Cell:

One of the more common and modern batteries is the nickel-cadmium or ni-cad cell.This is the cell

commonly used in:

o Video Cameras.

o Portable CD players.

o And any other cordless electrical device where you can recharge the battery from.The reactions also are reasonably simple:

Once again note to recharge these batteries the reactions are reversed.

Fuel Cells:

Fuel cells differ from primary cells in two major ways:

o STORE NEITHER REACTANTS OR PRODUCTS:

o The reactants are fed in at a constant rate.

o They convert this to supply electricity at a constant rate.

o REACTANTS ARE USUALLY GASEOUS:

o This is unlike the solid substances that you will usually find in primary or

secondary cells.

The most successful fuel cell so far is that based on the combination of oxygen and hydrogen:

o Hydrogen gas is supplied to the anode chamber.

o Oxygen gas is supplied to the cathode chamber

o The gasses then diffuse through the electrodes:

o These are usually porus metals such as platinum or nickel.

Reduction (Cathode): PbO2(s) + 4H+

(aq) +SO42!

(aq) + 2e!

"PbSO4( s) + 2H2O(l)

Oxidation (Anode): Pb(s) + SO42!

(aq) "PbSO4( s) + 2e!

Redox (Overall): Pb(s) +PbO2( s) + 4H+

(aq) + 2SO42!

(aq) " 2PbSO4( s) + 2H2O(l)

2PbSO4( s)+

2H2O(l)!

Pb(s)+

PbO2( s)+

4H

+

(aq)+

2SO42"

(aq)

Anode: Cd(s)+ 2OH!

(aq) " Cd OH( )2( s)+ 2e!

Cathode: NiO OH( )(s)

+ H2O(l)+ e!

"Ni OH( )2( s)

+ OH!(aq)

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of61 117o They act as catalysts.

o The gasses then react with the electrolyte, which may be acidic or alkaline.

This set up can be shown below:

Now if the electrolyte was acidic.

If the electrolyte was alkaline:

It is important to note, whether the electrolyte is acidic or alkaline we get the same combined

reaction:

This cell has a theoretical e.m.f of 1.23 Volts however this is difficult to attain:

o The H+and OH-ions (whether it is acidic or alkaline) do not migrate quickly between

electrodes, thus being a limiting factor on the speed of reaction.

o The electrodes do not have a great amount of contact with the gaseous molecules some

inventions including pores to help interaction has partially solved these problems.

Fuel cells are still part of a developing technology:

o They are relatively small.

o High fuel efficiency.

o This is why they are used in spacecraft.

And for this course that is as far as we need to go in the realm of fuel cells.

Electrolytic Cells & Electrolysis:

What we know so far is that ElectroCHEMICAL cells convert chemical energy into electrical

energy through redox reactions:

o Dry cells.

Anode: H2(g) ! 2H+

(aq) + 2e"

Cathode: O2(g) + 4H+

(aq) + 4e"

! 2H2O(l)

Anode: H2(g) + 2OH!

" 2H2O(l) + 2e!

Cathode: O2(g) + 2H2O(l) + 4e!

" 4OH!(aq)

2H2(g) +O2(g)! 2H2O(l)

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of62 117o Batteries et.

Basically there is an electron transfer (current) and to varying abilities the chemical energy produces

electricity.

AN ELECTROLYTIC CELL HOWEVER PASSES AN ELECTRIC CURRENT THROUGH A

SOLUTION AND TURNS ELECTRICAL ENERGY INTO CHEMICAL ENERGY!!!

Electrolysis Introduction:Basically we can say:

The Process of using an electric current to produce a chemical change is called ELECTROLYSIS

The important point to note here:

BECAUSE THE DIRECTION OF THE CURRENT IS REVERSED THE ELECTRICAL

CHARGE WILL BE REVERSED

So because of that we say:

o CATHODE:

o Electrode where reduction occurs.

! WILL BE NEGATIVE.

o ANODE:o Electrode where oxidation occurs.

! WILL BE POSITIVE.

There are still certain points that remain similar:

o Current Outside Cell caused by flow of electrons.

o Current inside Cell caused by movement of ions.

It is also worth noting:

o The E0for an ELECTROLYTIC CELL is NEGATIVE.

The only other point that is likely to come up is rather important:

ELECTROCHEMICAL CELLS EXOTHERMIC

ELECTROLYTIC CELLS - ENDOTHERMICThats all there really is to it some important features about ELECTROLYTIC CELLS (where

electrolysis occurs) are:

o Do not need to separate half cells because the reactionwill not occur spontaneously:

o However often the products of the electrolysis CAN recombine, in which case you DO

need to separate them.

o Power Supplied the electrodes are connected to an external power source (battery) so

they are charged.

Now we can safely move on to look at some examples of Electrolysis.

Electrolysis: Molten Salts:

WHENEVER you melt and ionic salt IT SPLITS UP INTO IONS:

By Josh Walker

7/21/2019 Chemistry

chemistry for first year university

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