Transcript H - Deans Community High School

Chemical Changes and
Structure
This unit covers (you get notes for this):
• How chemist can control the rate of
chemical reactions, and the enthalpy
changes that take place.
• Trends in the periodic table
• The relationship between the arrangement
of elements in the periodic table and their
bonding, structure and properties.
• Polar covalent bonds in the context of the
bonding
continuum,
followed
before
studying intermolecular forces.
Chemical changes and
Structure
From previous work you should know and understand
the following:
• Factors that affect the rate of a reaction and
average rate calculations.
• Atomic structure
• Electron orbital's or energy levels
• Valency
•Covalent and ionic bonding
•Physical properties of substances.
Controlling the rate
Overview
Learn to explain how a number
of key factors can influence
reaction rate, using the collision
theory.
Factors Affecting Rate of
Reaction
4 factors affect the speed of a reaction:
1. Particle size
2. Concentration of solution
3. Temperature
4. Catalysts
Collision Theory
a) Collision theory
Learning intention
Learn how chemists control
reaction rates by careful
consideration of the influence
of concentration, temperature,
surface area and collision
geometry.
Rates of Reaction
Reactions happen at different rates.
Industry needs to control reaction rates to
increase production and get a good return
for the investment
Rates may need to be controlled for
safety, or to keep the rate of production
within the limit of the plant
Collision Theory
For a chemical reaction to
occur, reactant molecules
must collide.
The collision must provide enough
energy to break the bonds in the
reactant molecules
Then new chemical bonds form to make product
molecules.
Progress of a Reaction
Reactions can be followed by measuring changes in concentration,
mass and volume of reactants and products.
A. Where is the reaction the quickest?
B. Why does the graph level off?
Rate
No more products formed.
C. Why does the graph curve?
A
The concentration of the
reactants decrease with time.
time
1.1 Effect of Surface Area
Effect of Surface Area
Particle size, the smaller the particles, the greater the
surface area, the greater the chance of successful collisions.
4X4= 16 cm2
16x6=96 cm2
2x2 = 4 cm2
4X6= 24 cm2
24X8= 192 cm2
Rate and Particle Size
Only the particles on the surface of a solid can
be involved in a collision
Higher Chemistry Eric Alan and John
Harris
Crushing a solid increases the surface area
more particles are available for collision
therefore increased rate of the reaction
Effect of Surface Area
Hydrochloric acid reacts with marble chips (calcium carbonate)
2HCl(aq) + CaCO3(s)
CaCl2(aq) + CO2(g) + H2O (l)
Rates of reaction
The rate of reaction can be followed by
measuring changes in
Concentration
Mass
Volume of gas produced
Measuring reaction rates
Change in
mass (g)
Products
Reactants
time (s)
Average rate
of reaction =
change in mass of product or reactant
time interval
Units
g s-1
Measuring reaction rates
Change in
volume (cm3)
Products
Reactants
time (s)
Average rate change in volume of product or reactant
of reaction =
in time for the change to occur
Units cm3 s-1
Measuring reaction rates
Change in
concentration (mol l-1)
Products
Reactants
Time (s)
Average rate change in concentration of product or reactant
of reaction =
time interval
Units mol l-1
s-1
How can we follow the
reaction?
A gas is produced. What will happen
to the gas if there is no lid on the
container?
What will happen to the mass?
How can we follow the rate?
What to do:
You can follow the rate of a reaction
involving gases by;
1.Measuring the volume of gases
produced over time
2. Measuring the loss of mass over
time
How can we follow the
reaction?
• If we use a container fitted with a delivery
tube we could measure the amount of gas
produced. How?
Measuring rate of reaction
Two common ways:
1) Measure how fast the
products are formed.
•
method can accurately
measure volume of gas
produced.
2) Measure how fast the
reactants are used up by
measuring the mass lost.
Following the rate of a
reaction–Activity 1.1
Follow the instructions on the experiment card
Write down the aim, next to your note.
Complete the experiment , recording your
results and write a conclusion – complete all
questions on the back of the instructions.
What to do
Record your results in a the table.
Plot a graph of volume vs time using the same
axes for both sets of data
rate = change in volume ( the unit is cm3 s-1)
time interval
Swap results
Each group should have a sets of
results which can be used to plot
graphs.
2.5
2
Loss in mass (g)
1.5
lumps
powder
1
0.5
0
0
20
40
60
80
100
Time (s)
rate = change in mass = ____________________ g s-1
time interval
Rate over 1st 25 seconds
(g s-1)
rate over 2nd 25seconds
(g s-1)
Whole chips (C)
0.8-2
25-0
=0.05
0.35 -0.8
50-25
=0.018
Ground chips (G)
0.3-2
25-0
=0.068
0.25-0.3
50-25
=1x10-3
Now do Question 1-5 on back of instructions.
Assessment Activity 1.1 Marble Chips and Acid
1.
This experiment used changes in mass to follow the rate of the
reaction. Suggest other methods of monitoring a chemical reaction
and thus determining its rate.
2. Write a balanced equation for the reaction between marble chips
(calcium carbonate) and dilute hydrochloric acid.
3. Describe the way the course of this reaction was followed.
4. Why does the curve level off after some time?
5.
From your small chips graph, estimate the time it takes for
- all the acid to react,
- half the acid to react.
6. Explain why the time taken for all the acid to react is not exactly twice
the time taken for half the acid to react.
Effect of Concentration
Rate and Concentration
For a reaction to
take place the
particles must
collide.
Increasing the concentration of a solution
increases the number of particles in the same
volume.
Therefore more chance
of collision i.e.
increased rate of the reaction
Effect of Concentration
• The higher the concentration, the more
particles in a given space, the more chance
there is of successful collisions.
Using Relative (relevant) Rate
1. A reaction with high concentration took 90 seconds to complete.
A similar reaction with low concentration took 160 seconds to
complete. Calculate the relative rate of both reactions.
2. a) Calculate the time for reaction 1 when concentration is
0.4mol/l.
b) Calculate the time for reaction 2 at 40oC
Effect of Concentration –the chemical
clock challenge
Your challenge is to create a series of
solutions that will change colour in
time to music.
http://www.youtube.com/watch?v=rSAa
iYKF0cs
Effect of Concentration –the chemical
clock challenge
Checklist;
•Aim of the experiment
•Hypothesis
•Method, which variables to control and change
•What to measure and how to measure it
•How to record your results – table of results
•What graph to draw – see page 6 of notes
•Make a conclusion
•Evaluate – how can you improve it/reduce the
errors
Effect of Concentration –the chemical
clock challenge
The reaction is between potassium
iodide (KI) and hydrogen peroxide.
The iodine clock reaction changes from
colourless to blue/black
Effect of Concentration –the chemical
clock challenge
You will carry out the reaction using a series of
dilutions of the iodide solution. This will be
diluted by replacing some of the volume with
water.
Effect of concentration –the chemical clock
challenge
2I- (aq) + H2O2
(aq)
I2 (aq) + 2S2O32-
+ 2H+
(aa)
(aq)
 2H2O
(l)
+ I2
 2I- (aq) + S4O62-
(aq)
(ag)
The reaction mixture stays colourless as the iodine molecules are
converted back to iodide molecules by the thiosulphate ions.
Once all the thiosulphate ions have been used, a blue black colour appears
suddenly as iodine reacts with starch.
Relative Rate =
t is the time….
1
t
Units s-1
This is a measure of how long it takes
for the blue/black colour to form. (when excess I2 forms)
Effect of concentration –the chemical clock
challenge
1) Using syringes measure out
10cm3 sulphuric acid 0.1moll-1
10cm3 sodium thiosulphate 0.005moll-1
1cm3 starch solution
25cm3 potassium iodide solution 0.1mol l-1
Into a dry 100cm3 beaker
2) Measure out 5cm3 of hydrogen peroxide 0.1moll-1
into a syringe. Add it to the mixture as quickly as possible and
start the timer.
3) Stop the clock when the mixture suddenly turns dark blue.
4) Repeat, using 20 cm3 of potassium iodide solution and 5cm3
of water with, then using repeated dilutions
Effect of Concentration –the chemical clock
challenge
Volume Volume Concentra
of water of 0.5 tion of KI
(cm3)
mol l-1
KI (aq)
(cm3)
O.0
25.0
0.1
5.0
20.00
10.0
15.0
15.0
10.0
20.0
5.0
0.06
0.02
1st
Time
(s)
2nd
time
Average Relative
time
Rate
(1/t)
Effect of Concentration –the chemical
clock challenge
RESULTS - Plot a graph showing the
volume of potassium iodide x axis and the
rate of reaction on the y axis.
Effect of Changing Concerntration on Rate of Reaction
Write-up
1.
2.
3.
4.
5.
6.
Aim
Hypothesis
The independent variable (what you changed),
The dependent variable (what you measured),
Safety
The method mentioning all the equipment used and measurements
made, readings and variable kept constant/changed etc
7. A table (with headings) of your measurements, and a sample
average and rate = 1/t calculation) and your line graph.
8. Your conclusion (what you found out – must mention results and
link to the aim)
9. An evaluation (Comment on how well you felt exp. Went and how
you can improve your investigation)
Effect of Concentration - Lemonade
challenge
• Your task is to make
up a solution
containing the same
concentration of
sugar as a can of
lemonade which
contains 24g of
sucrose
C12 H22 O11
Effect of Concentration - Lemonade
challenge
Work out the mass of sucrose
required to make up 100cm3 of
sucrose solution of the same
concentration as lemonade, assuming
there are 24g in 330cm3
Make up your solution.
What is the concentration of this
solution?
Making a standard sample
• The sample is weighed and dissolved in a small
volume of (deionised) water in a beaker.
• The solution is transferred to a standard flask.
• The beaker is rinsed and the rinsings also poured
into the standard flask.
• The flask is made up to the mark adding the last
few drops of water using a dropping pipette.
• The flask is stoppered and inverted several times
to ensure thorough mixing of the solution.
Effect of Concentration - Lemonade
challenge
Concentration (c ) is measured in
moles per litre ( mol l-1)
no moles =
cxv
1000
To calculate the concentration you need
to work out the number of mol of
sugar present.
Effect of Concentration - Lemonade
challenge
No moles = mass (g)
GFM
GFM = gram formula mass
Effect of Concentration - Lemonade
challenge
• Sucrose is a non-reducing sugar – it does
not react with Benedicts unless it is first
hydrolysed.
• Boil 10 cm3 sugar solution with 5cm3 1mol/l
HCl. Then neutralise the solution with 5cm3
1mol/l NaOH.
• Allow this solution to cool to room
temperature.
• Repeat with the lemonade solution.
• Add 2cm3 of Benedicts to each sample.
• Prepare a beaker of boiling water.
Effect of Concentration - Lemonade
challenge
Add the test tubes to the boiling
water
If you have made up the solutions
correctly, all your solutions should
take the same time to change colour.
Effect of Temperature
on Rate of Reaction
Effect of Temperature -the vanishing cross
Sodium thiosulfate solution is reacted with acid.
A precipitate of sulfur forms. The time taken
for a certain amount of sulfur to form is used
to indicate the rate of the reaction.
Effect of Temperature -the vanishing cross
Effect of Temperature -the vanishing cross
Effect of Temperature -the vanishing cross
Results
Temperature
(0C)
Reaction
time in
seconds
1/time
19
32
38
51
60
105
46
36
18
12
0.0095
0.0217
0.0278
0.0556
0.0833
Effect of Temperature on Rate of Reaction
0.1
0.09
A 10 oC rise in
temperature will
approximately double the
rate of the reaction
0.08
0.07
1/Time
0.06
0.05
0.04
0.03
0.02
0.01
0
0
10
20
30
Temperature
40
50
60
70
Temperature and Energy
Effect of Temperature -the vanishing cross
• The experiment can be viewed at
• http://media.rsc.org/Classic%20Chem%20e
xperiments/CCE-64.pdf
Activation energy and reaction pathway
Products
Reactants
1.
Before collision
Reactants
Energy changes
2. In full collision
Activated complex
Partially broken reactant
bonds and partially
formed product bonds
3. After collision
Products
If the reactants have enough combined K.E. to
overcome Ea , their K.E. is converted into the energy
needed to form the activation complex.
Temperature and kinetic energy
What do we mean by temperature and heat (thermal energy) ?
The thermal energy of a system is a measure
of both the potential and kinetic energy within
the system.
The temperature is a measure of how ‘hot’
a system is.
The temperature is a measure of the average kinetic energy in
a system.
Temperature and Energy
Energy Distribution
No of
molecules
Ea
Number of collisions which
result in new products
being formed.
Kinetic energy
Total number of collisions with sufficient K.E. energy is
the area under the graph to the right of the Ea .
No of
molecules
Temperature and Activation energy
Increasing the temperature means a greater
Activation
EA
number of molecules have energies
in excessenergy
of
At the higher temperature
EA.
greyof
area
T2 theThe
number
particles
T1
T2
represents
the of E
with energy
in excess
A
number
particles
is greater.
Theofblue
area now
witha energy
represents
particles
with
Even a small rise in temperature causes
largein excess
ofinthe
EA. These
energy
excess
of EA.
increase in the number of particles
with
energy
particles have enough
above EA. Therefore a greater proportion
energyof
to cause
successful collisions
collisions will be successful.
Energy associated with
molecules
The blue area is larger than the
grey so at temperature T2 more
particles have enough energy to
cause successful collisions.
Energy Distribution-Concentration
No of
molecules
Ea
Ea does not change but the number of successful
collisions increases significantly, so rate increases.
Increased temperature
Kinetic energy
No of
molecules
Ea
Increased concentration
Ea does not change but the number of successful
collisions increases.
Kinetic energy
b) Reaction Profiles
Learning intention
• Learn how a potential energy
diagram can be used to describe
a reaction pathway, and to
display activation energy and
reaction enthalpy.
• Learn
that
the
activated
complex
is
the
unstable
intermediate formed at the peak
of the potential energy diagram.
Reaction Profiles – Enthalpy (clip)
Thermochemistry is the study of heat energy taken in
or given out in chemical reactions. This heat, absorbed
or released, can be related to the internal energy of
the substances involved. Such internal energy is called
ENTHALPY, symbol H.
As it is only possible to measure the change in enthalpy,
the symbol  H, is used.
 H = Hp - Hr
Enthalpy (products) – Enthalpy (reactants)
Units kJ, kilojoules
Reaction profiles -Exothermic and Endothermic Reactions
Step 1: Energy must be
SUPPLIED to break bonds:
Step 2: Energy is RELEASED
when new bonds are made:
A reaction is EXOTHERMIC if more energy is RELEASED
then SUPPLIED. If more energy is SUPPLIED then is
RELEASED then the reaction is ENDOTHERMIC
Reaction profiles -Exothermic and Endothermic Reactions
-H
Enthalpy of
reactants
Enthalpy of
products
+H
Enthalpy of
reactants
Enthalpy of
products
Exothermic reactions give out thermal energy so the enthalpy of the
products is less than that of the reactants.
Endothermic reactions take in thermal energy from their surroundings.
The enthaply of the products is greater than that of the
reactants
Reaction profiles -Exothermic and Endothermic Reactions
Exothermic reactions give out heat,
causing a rise in the temperature
Endothermic reactions take in heat
The energy change in a reaction can be
shown in a potential energy diagram
or reaction profile
-H
Path of reaction
Exothermic reactions
give out thermal energy
 H = -ve
PE kJmol-1
PE kJmol-1
Reaction profiles -Exothermic and Endothermic Reactions
+H
Path of reaction
Endothermic reactions take in
thermal energy from their
surroundings.
 H = +ve
Reaction profiles -Exothermic and Endothermic Reactions
A. Combustion of methane
CH4
H
(g)
+ 2O2
CH4
CO2
(g)
+ 2 H20
products
B. Cracking of ethane
C2H4
(g) +
products
H
kJmol-1
C2H6
(g)
reactants
+ 2O2
(g)
→ CO2
(g)
+ 2 H20
(l)
(g)
 H negative, exothermic reaction
reactants
kJmol-1
(g)
C2H6
(l)
(g)
→ C2H4
(g)
+ H2(g)
H2(g)
 H positive, endothermic reaction
Reaction profiles -Exothermic and Endothermic Reactions
Mix the following pairs of chemicals in a
polystyrene cup to discover if the
reactions are exothermic or endothermic
Reaction
10cm3 NaOH
+ 10cm3 HCl
10cm3
NaHCO3 + 4
spatulas
citric acid
10cm3 CuSO4
+ spatula of
Zn powder
10cm3 H2SO4
+ Mg ribbon
Temp
before
mixing/oC
Temp after Endothermic
mixing/oC
or
exothermic
Use of the thermite reaction
This reaction is used to weld railway
lines together
http://www.youtube.com/watch?v=vCqG
3rWtNbc&feature=related
An endothermic reaction
Higher Chemistry Eric Alan and John Harris
Enthalpy changes and industrial
processes
For industrial processes it is essential
that chemists can predict the
quantity of heat taken in or given out.
Exothermic reactions lower the
temperature, slowing the reaction
rate
Heat must be supplied to maintain the
rate of reaction – this is an expense
Enthalpy changes and industrial
processes
Exothermic processes produce heat
Heat may need to be removed to
prevent the reactions proceeding
beyond the capacity of the plant
c) Temperature and kinetic
energy
Learning intention
Learn to describe the
relationship between
temperature, kinetic energy and
activation energy.
Activation energy and reaction pathway
Potential energy diagrams give useful information
about the energy profile of a reaction.
The activation energy is the minimum kinetic
energy required by colliding molecules for a
reaction to occur. In the diagrams shown above
the activation energy appears like a ‘energy
barrier’ which reactants must get over to become
products.
Higher Chemistry Eric Alan and John Harris
Activation energy and reaction pathway
2
Breaking bonds
1
Making bonds
P.E.
3
Activation Energy EA
1
2
Activated complex
3
Reaction Path
Activation Energy is the additional P.E. which has to be
attained by colliding molecules to form an activated complex.
Activated complex is the unstable arrangement of atoms
formed at the maximum of the potential energy barrier.
Activation energy and reaction pathway
As a reaction proceeds from reactants to products, an
intermediate stage is reached at the top of the activation
barrier at which a highly energetic species called an activated
complex is formed.
A+B
→
X
→
C+D
Higher Chemistry Eric Alan and John Harris
Activation energy and reaction pathway
A+B
→
X
→
C+D
This unstable activated complex only exist for a short period
of time. From the peak of the energy barrier it can lose
energy in one of two ways i.e. to the stable products or to
form the reactants again.
The higher the Ea the higher the barrier and the slower
the reaction.
Higher Chemistry Eric Alan and John Harris
Ea
-H
Exothermic reactions
give out thermal energy
 H = -ve
Enthalpy
Enthalpy
Activation energy and reaction pathway
Ea
+H
Endothermic reactions take in
thermal energy from their
surroundings.
 H = +ve
Activation energy and reaction
pathway
1. Mark Ea and ∆H on the PE diagrams and then calculate the value of
each for the forward reaction.
Ea
Ea
Ea
∆H
∆H
∆H
A Ea = 50 KJmol-1
∆H = -10 kJmol-1
B Ea = 30 KJmol-1
C Ea = 40 KJmol-1
∆H = -40 kJmol-1
∆H = +20 kJmol-1
Higher Chemistry Eric Alan and John Harris
Activation energy and reaction
pathway
2. Mark Ea and ∆H on the PE diagrams and then calculate the value of
each for the reverse reaction.
Ea
Ea
Ea
∆H
∆H
∆H
A Ea = 60 KJmol-1
B Ea = 70 KJmol-1
C Ea = 20 KJmol-1
∆H = +10 kJmol-1
∆H = +40 kJmol-1
∆H = -20 kJmol-1
Higher Chemistry Eric Alan and John Harris
Activation energy and reaction pathway
Ea
Calculate Ea for the forward reaction
Ea = 210 – 20 = 190kJ
Higher Chemistry Eric Alan and John Harris
d) Catalysts
Learning intention
Learn how a catalyst speeds up
reaction rate by lowering the
activation energy, and how to
represent this on a potential
energy diagram.
Catalysts at Work
Heterogeneous
When the catalyst and reactants are in different
States you have ‘Heterogeneous Catalysis’.
They work by the adsorption of reactant molecules.
E.g. Ostwald Process (Pt) for making nitric acid and
the Haber Process (Fe) for making ammonia and
the Contact Process (Pt) for making Sulphuric Acid.
Homogeneous
When the catalyst and reactants are in the same state
you have ‘Homogeneous Catalysis’. E.g. making ethanoic acid
from methanol and CO using a soluble iridium complex.
Enzymes
Are biological catalysts, and are protein molecules
that work by homogeneous catalysis. E.g. invertase
and lactase.
Enzymes are used in many industrial processes
How a heterogenous catalyst works
Heterogenous Catalysis are thought to work in three
stages...
Adsorption
Reaction
Desorption
Higher Chemistry Eric Alan and John Harris
How a heterogenous catalyst works
For an explanation of what happens click on the numbers in turn, starting with

How a heterogenous catalyst works
Adsorption (STEP 1)
Incoming species lands on an active site and forms bonds with the catalyst. It may use some of
the bonding electrons in the molecules thus weakening them and making a subsequent reaction
easier.
How a heterogenous catalyst works
Adsorption (STEP 1)
Incoming species lands on an active site and forms bonds with the catalyst. It may use some of the
bonding electrons in the molecules thus weakening them and making a subsequent reaction easier.
Reaction (STEPS 2 and 3)
Adsorbed gases may be held on the surface in just the right orientation for a reaction to occur.
This increases the chances of favourable collisions taking place.
How a heterogenous catalyst works
Adsorption (STEP 1)
Incoming species lands on an active site and forms bonds with the catalyst. It may use some of
the bonding electrons in the molecules thus weakening them and making a subsequent reaction
easier.
Reaction (STEPS 2 and 3)
Adsorbed gases may be held on the surface in just the right orientation for a reaction to occur.
This increases the chances of favourable collisions taking place.
Desorption (STEP 4)
There is a re-arrangement of electrons and the products are then released from the active sites
Examples of heterogenous catalysts
Format
Metals
Ni, Pt
Fe
Rh, Pd
hydrogenation reactions
Haber Process
catalytic converters
Oxides
Al2O3
V2O5
dehydration reactions
Contact Process
FINELY DIVIDED
increases the surface area
provides more collision sites
IN A SUPPORT MEDIUM
maximises surface area and reduces costs
How a heterogenous catalyst works
In some cases the choice of catalyst can influence the products
Ethanol undergoes different reactions depending on the metal used
as the catalyst.
The distance between active sites and their similarity with the
length of bonds
determines the method of adsorption and affects which bonds are
weakened.
Copper
Dehydrogenation (oxidation)
C2H5OH ——>
CH3CHO + H2
Alumina Dehydration
C2H5OH ——>
C2H4 + H2O
How a heterogenous catalyst works
Poisoning
Impurities in a reaction mixture can also adsorb onto the surface
of a catalyst thus removing potential sites for gas molecules and
decreasing efficiency.
expensive
because...
the catalyst has to be replaced
the process has to be shut
down
examples
Sulphur poisons iron in Haber process
Lead poisons Pt in catalytic converters in
cars
Investigating Catalysis
Decomposition of hydrogen peroxide
Hydrogen → water + oxygen
2H2O2
→ 2H2O + O2
What will you see during the reaction?
What is the test for oxygen?
Which type of catalysis?
Did the experiment involve heterogeneous or
homogeneous catalysis?
An Example of a homogenous catalyst
Higher Chemistry Eric Alan and John Harris
Catalysts and Potential energy diagrams
Catalysts work by providing…
“AN ALTERNATIVE REACTION PATHWAY WHICH HAS
A LOWER ACTIVATION ENERGY”
Potential energy graphs and catalysts
Uncatalysed
75 -
60 50 -
Catalysed
Reactants
P.E.
Products
25 -
Reaction path
Potential energy graphs and catalysts
Activation energy Ea for the forward uncatalysed reaction
75 -
60 -
Activation energy Eafor the
forward catalysed reaction
Reactants
50 -
P.E.
Products
25 -
Reaction path
Catalysts lower the activation energy needed for a successful collision.
Potential energy graphs and catalysts
Activation energy Ea for the reverse uncatalysed reaction reaction
75 -
Activation energy Ea for the
reverse catalysed reaction
60 50 -
Reactants
P.E.
Products
25 -
Reaction path
Catalysts lower the activation energy needed for a successful collision.
Potential energy graphs and catalysts
75 -
Catalysts lower the activation
energy needed for a successful
collision.
60 50 -
Reactants
P.E.
Products
25 -
Reaction path
∆H
Effect of catalyst – forward reaction
Effect of catalyst – reverse reaction
Activation energy
No change Lowered
No change Lowered
Catalysts
A catalyst speeds up the reaction by lowering the
activation energy.
A catalyst does not effect the enthalpy change for a
reaction
A catalyst speeds up the reaction in both directions
and therefore does not alter the position of
equilibrium or the yield of product, but does decrease
the time taken to reach equilibrium.
Energy distribution and catalysts
Ea
No of
Collisions
with a
given
K.E.
Un-catalysed reaction
Kinetic energy
Ea
Total number of collisions (area under
the graph) with sufficient K.E.
energy to create new products.
Catalysed reaction
Ea is reduced
Concentrated solutions of hydrogen peroxide are used in the
propulsion systems of torpedoes. Hydrogen peroxide
decomposes naturally to form water and oxygen:
2H2O2(aq) → 2H2O(ℓ) + O2(g)
ΔH = −196∙4 kJ mol–1
Transition metal oxides act as catalysts in the decomposition of
the hydrogen peroxide.
Unfortunately, there are hazards associated with the use of
hydrogen peroxide as a fuel in torpedoes. It is possible that a
leak of hydrogen peroxide solution from a rusty torpedo may
trigger an explosion.
Using your knowledge of chemistry, comment on why this could
happen.
Trends in the periodic
table and bonding
Overview
This section studies
• how the elements are arranged in
the Periodic Table,
• the structure and bonding of the
first twenty elements,
• Helps you explain key periodic
trends in physical properties and
relate these to the bonding
continuum.
a) The arrangement of
elements in the periodic table
Learning intention
• Learn how the elements are
organised into groups and
periods in order of increasing
atomic number
• To
identify
important
classifications of elements
within the periodic table.
It has taken many years of work by many
scientists to find out about the elements that we
know about now (and there may be more that we
don’t know about yet).
Ancient Times - Prior to 1
A.D.
Gold
Silver
Copper
Iron
Mercury
Carbon
Sulphur
LeadTin
Timeline of the elements
Arsenic (Magnus ~1250)
Antimony (17th century or earlier)
Phosphorus (Brand 1669)
Zinc (13th Century India)
Cobalt (Brandt ~1735)
Platinum (Ulloa 1735)
Nickel (Cronstedt 1751)
Bismuth (Geoffroy 1753)
Robert Boyle
In 1661, Robert Boyle showed
that there were more than just
four elements as the ancients
had assumed.
Boyle defined an element as a
pure substance that cannot be
decomposed into any simpler
substance.
Time line of elements
Hydrogen (Cavendish 1766)
Nitrogen (Rutherford 1772)
Oxygen (Priestley; Scheele 1774)
Chlorine (Scheele 1774)
Manganese (Gahn, Scheele, & Bergman 1774)
Molybdenum (Scheele 1778)
Tungsten (J. and F. d'Elhuyar 1783)
Tellurium (von Reichenstein 1782)
Lavoisier 1789
http://web.bilkent.edu.tr/
The first modern list of chemical elements was
given in Antoine Lavoisier's 1789 Elements of
Chemistry, which contained thirty-three
elements, including light.
Introduced a logical system for naming compounds
and helped introduce the metric system
Time of the chemists
Uranium (Peligot 1841)
Strontium (Davey 1808)
Titanium (Gregor 1791)
Yttrium (Gadolin 1794)
Vanadium (del Rio 1801)
Chromium (Vauquelin 1797)
Beryllium (Vauquelin 1798)
Niobium (Hatchett 1801)
Tantalum (Ekeberg 1802)
Atomic Weights
John Dalton, 1803, was the first chemist to use the term ‘atom’
He used this idea to explain how elements react together to
form molecules.
Dalton suggested that it should be possible to compare the
masses of atoms.
Hydrogen
1
Carbone
4.2
Oxygen
5.5
Water
6.5
Sulphur
14.4
Sulphuric Acid
25.4
www.bioanalytical.com
Cerium (Berzelius & Hisinger;
Klaproth 1803)
Palladium (Wollaston 1803)
Rhodium (Wollaston 18031804)
Osmium (Tennant 1803)
Iridium (Tennant 1803)
Sodium (Davy 1807)
Potassium (Davy 1807)
Barium (Davy 1808)
Calcium (Davy 1808)
Magnesium (Black 1775; Davy
1808)
Boron (Davy; Gay-Lussac &
Thenard 1808)
Iodine (Courtois 1811)
Lithium (Arfvedson 1817)
Cadmium (Stromeyer 1817)
Selenium (Berzelius 1817)
Silicon (Berzelius 1824)
Zirconium (Klaproth 1789;
Berzelius 1824)
Aluminum (Wohler 1827)
Bromine (Balard 1826)
Thorium (Berzelius 1828)
Lanthanum (Mosander 1839)
Terbium (Mosander 1843)
Erbium (Mosander 1842 or
1843)
Ruthenium (Klaus 1844)
Cesium (Bunsen & Kirchoff
1860)
Rubidium (Bunsen & Kirchoff
1861)
Thallium (Crookes 1861)
Indium (Riech & Richter 1863
Newlands
Newlands in 1865, placed elements in order of
www.chemsoc.org
succession of atomic weights noticed a pattern, noticed
that the 8th one was a ‘kind of repetition of the 1st.
He called this the ‘Law of Octaves’.
Element
Atomic
weights
Element
Atomic
Weights
Element
Atomic
Weights
Hydrogen
1
Fluorine
8
Chlorine
15
Lithium
2
Sodium
9
Potassium
16
Beryllium
3
Magnesium
10
Calcium
17
Boron
4
Aluminium
11
Chromium
18
Carbon
5
Silicon
12
Titanium
19
Nitrogen
6
Phosphorus
13
Manganese
20
Oxygen
7
Sulphur
14
Iron
21
Lothar Meyer
www.apsidium.com
Meyer in 1869, independently, put forward a similar list of
elements.
Meyer plotted graphs of
melting point, boiling point
and atomic volume against
atomic mass.
He found the properties
varied in a regular way i.e. periodically
Mendeleev (1869)
In 1869 he published ‘Principles of Chemistry’
- proposed the modern Periodic Table.
elmoscow.ru
- elements with similar properties were placed
together
- he left gaps for new 'undiscovered' elements.
- he predicted properties of undiscovered elements
- arranged in order of increasing relative atomic mass
- some elements were out of order therefore modern
table is arranged in Atomic Number
Meyer recognised Mendeleev’s work and both where
awarded The Davy medal for Chemistry in 1882.
The world’s first view of Mendeleev’s Periodic Table
– an extract from Zeitschrift fϋr Chemie, 1869.
Correct predictions
The greatness of Mendeleev was that not only did
he leave spaces for elements that were not yet
discovered but he predicted properties of five of
these elements and their compounds including
gallium which he called eka-aluminium.
In Paris (1875) Paul Emile Lecoq de Boisbaudran discovered an element he
named gallium after the Latin name for France
Eka-aluminium (Ea)
Gallium (Ga)
Atomic weight
About 68
69.72
Density of solid
6.0 g/cm3
5.9 g/cm3
Melting point
low
29.78oC
Valency
3
3
Method of discovery
Probably from its
spectrum
Spectroscopically
Oxide
Formula Ea2O3, density
5.5 g/cm3. Soluble in
both acids and alkalis.
Formula Ga2O3, density
5.88 g/cm3. Soluble in
both acids and alkalis.
Sir William Ramsay
One thing that Mendeleev did not predict was the
discovery of a whole new Group of elements, the
noble gases, discovered by Scot William Ramsay
and
th
co-workers during the last decade of the 19
century.
Groups
The Periodic Table
Periods
Groups
- vertical columns.
- elements in a group
show specific
similarities.
- common names : Alkali
Metals, Halogens,
Noble Gases.
- increasingly metallic
down a group, nonmetallic up a group.
- outer shell electrons
determine the group
number.
Periods
- horizontal rows
- two short periods, four
long periods.
- elements change from
metallic to nonmetallic across a
period.
- number of the shell
determines the period.
Periodicity
It is the ideas of Meyer and Mendeleev that we will
make use of to try and understand the relationships
between the first 20 elements.
Periodicity -
As you move across a Period from left to right, a pattern
emerges. A similar pattern appears on crossing the next
Period.
•Properties which behave in this way are said to be periodic.
Periodicity is the regular recurrence of similar element
properties.
•
Density - Lothar Meyer
Curve
Variation of density (g cm-3) with atomic number
Adapted from New Higher Chemistry E Allan J Harris
period 2 (Li - Ne)
maximum at boron (B) group3
period 3 (Na - Ar) maximum at
Aluminium (Al)- group 3
Variation of density (g cm-3) with atomic number
Adapted from New Higher Chemistry E Allan J Harris
In general in any period of the table, density first increases from group 1 to
a maximum in the centre of the period, and then decreases again towards
group 0
5th
4th
2nd
3rd
Variation of density (g cm-3) with atomic number
Adapted from New Higher Chemistry E Allan J Harris
down a group gives an overall increase in
density
Melting point - Lothar Meyer
Curve
Variation of melting point with atomic number
Adapted from New Higher Chemistry E Allan J Harris
Determined by the strength of intermolecular
bonding, between particles
period 2, peak at carbon
period 3, peak at silicon
In general the forces of attraction
(intermolecular bonding) for elements
on the left of the table must be
stronger, or more extensive than
between the particles on the right.
Variation of melting point with atomic number
Adapted from New Higher Chemistry E Allan J Harris
Down group 1 the alkali metals m.pt.
decrease there must be a decrease in
the force of attraction between the
particles
Variation of melting point with atomic number
Adapted from New Higher Chemistry E Allan J Harris
Down group 7 the halogens m.pt.
increases there must be a increase in
the force of attraction between the
particles
Boiling point - Lothar Meyer
Curve
Variation of boiling point with atomic number
Adapted from New Higher Chemistry E Allan J Harris
period 2, peak at carbon
period 3, peak at silicon
In general we see the same trend in
boiling point across the period
Variation of boiling point with atomic number
Adapted from New Higher Chemistry E Allan J Harris
Down group 1 the alkali metals b.p.
decrease once again there must be a
decrease in the force of attraction
between the particles
Variation of boiling point with atomic number
Adapted from New Higher Chemistry E Allan J Harris
Down group 7 the halogens b.p. increases
once again there must be a increase in
the force of attraction between the
particles
b) Periodic trends in covalent
radii and ionisation energies
Learning intention
• Learn the definitions of
covalent radius and first
ionisation energy
• find out how to explain their
periodic trends in relation to
atomic size, nuclear charge
and the screening effect of
inner shells of electrons.
Covalent radius
Atomic Size
There is no definite
edge to an atom.
However, bond lengths
can be worked out.
Covalent radius,
½ the distance between nuclei.
To find the bond length, add 2 covalent radii together.
pm = picometre X 10
– 12
m
Variation of covalent radius with atomic number
Adapted from New Higher Chemistry E Allan J Harris
The covalent radii of the elements in
any period decrease with increasing atomic number.
Variation of covalent radius with atomic number
Adapted from New Higher Chemistry E Allan J Harris
The covalent radii of the elements in
any group increase with increasing atomic number.
Variation of covalent radius with atomic number
Adapted from New Higher Chemistry E Allan J Harris
No values are given for the Nobel gases
Why?
Unreactive so do not form
bonds
Covalent radius
COVALENT RADIUS
Explain the change in covalent radius as you go along a period.
The covalent radius of an element is half the distance
between the nuclei of two of its covalently bonded
atoms.
The covalent radius decreases as you go along a period
As you go along a period there is a greater positive
charge on the nucleus The shells or energy levels of
electrons are more strongly attracted to the nucleus
and therefore the size of the atoms decreases.
COVALENT RADIUS
Explain the change in covalent radius as you go down a group.
The covalent radius of an element is half the
distance between the nuclei of two of its
covalently bonded atoms.
The covalent radius increases as you go down a group.
As you go down a group there are more energy levels of
electrons. The outer electrons are further away from
the nucleus so the atoms are larger.
Trend in Ionisation energy
Ionisation energies
This is defined as "the amount of energy
required to remove one mole of electrons from
one mole of atoms in the gaseous state”
Energy
e e
M (g)  M+(g) + e 1st ionisation
+
M +(g)
The outermost electron will be
the most weakly held and is
removed first
Ionisation energies
This is defined as "the amount of energy required to
remove one mole of electrons from one mole of atoms in
the gaseous state”
Energy
e
M (g)  M+(g) + e 1st ionisation
e
+
M
2+
(g)
M(g)+  M(g)2+ + e 2nd ionisation
The ionisation energy is an enthalpy change and
therefore is measured per mole.
Units kJmol-1 (kilojoules per mole).
Ionisation energies kJ mol-1
Overall increase along period
Decrease down group
Li
526
Be
905
B
807
C
N
O
F
Ne
1090 1410 1320 1690 2090
Na
502
K
425
Rb
409
Mg
744
Ca
596
Sr
556
Al
584
Ga
577
In
556
Si
792
Ge
762
Sn
715
P
1020
As
953
Sb
816
S
1010
Se
941
Te
870
Cl
1260
Br
1150
I
1020
Ar
1530
Kr
1350
Xe
1170
FIRST IONISATION ENERGY
Explain the change in first ionisation energy as you go
down a group.
The first ionisation energy is the energy required to
remove 1 mole of electrons from 1 mole of atoms in
the gaseous state.
The first ionisation energy decreases as you go down
a group.
As you go down a group there are more energy levels
of electrons. The outer electron is further from the
nucleus. The inner electrons shield the outer
electron from the effect of the nucleus. Less
energy is needed to remove the outer electron.
First and Second ionisation energies of the first 20 elements
Adapted from New Higher Chemistry E Allan J Harris
Down a group first ionisation energy
decreases
FIRST IONISATION ENERGY
Explain the change in first ionisation energy as you go
along a period.
The first ionisation energy is the energy required
to remove 1 mole of electrons from 1 mole of
atoms in the gaseous state.
The first ionisation energy increases as you go
along a period
As you go along a period there is a greater
positive charge on the nucleus. There is a
greater attraction between the outer electron
and the nucleus. More energy needs to be
supplied to remove the electron.
First and Second ionisation energies of the first 20 elements
Adapted from New Higher Chemistry E Allan J Harris
In each period there is an overall
increase peaking at the noble gas
First and Second ionisation energies of the first 20 elements
Adapted from New Higher Chemistry E Allan J Harris
For each element the second ionisation energy is
higher than the first ionisation energy.
First and Second ionisation energies of the first 20 elements
Adapted from New Higher Chemistry E Allan J Harris
It is worth noting the Nobel gases have the highest value
for each period. This goes some way to explaining the
great stability of filled orbital's and the resistance of
the Nobel gases to form compounds.
Successive ionisation Energies
first ionisation energy
E(g)
second ionisation energy E+(g)
third ionisation energy E 2+(g)
fourth ionisation energy E 3+(g)
E+(g)
+ eE 2+ (g) + eE 3+ (g) + eE 4+ (g) + e-
- ionisation energies increase as successive electrons
are removed
- removing an electron from a filled inner shell
requires a large increase in energy
The first four ionisation energies of aluminium, for
example, are given by
Al(g)
Al+(g)
Al2+(g)
Al+(g)
+ e-
1st I.E. = 577 kJ mol-1
Al2+ (g) + e-
2nd I.E. = 1820 kJ mol-1
Al3+ (g)
3rd I.E. = 2740 kJ mol-1
+ e-
In order to form an Al3+(g) ion from Al(g) you would have to supply:
577 + 1820 + 2740 = 5137 kJ mol-1
Explain why the second ionisation energy of an element
is always greater than the first ionisation energy:
First ionisation energy –
first mole of electrons
removed
Second ionisation energy –
second mole of electrons
removed
M(g)  M+(g) + e
M+(g)  M2+(g) + e
In the second ionisation energy negative electrons
are being removed from positive ions rather than
neutral atoms.
In the positive ion there is a greater attraction for
the electrons so more energy is needed to remove
the second mole of electrons.
Explain why the second ionisation energy of K is much
greater than the second ionisation energy of Mg:
K
(g)
2,8,8,1
 K+ (g) + e
2,8,8
K+ (g)  K2+ (g) + e
2,8,8
2,8,7
Mg (g)  Mg+ (g) + e
2,8,2
2,8,1
Mg+ (g)  Mg2+ (g) + e
2,8,1
2,8
The second ionisation of K
involves removing an
electron from a stable
electron arrangement.
The second ionisation of
Mg involves removing an
electron to form a stable
electron arrangement.
This requires a lot of energy
This requires less energy
Nobel gas compounds
If some other change can compensate for the energy
required then ionic compounds of Nobel gases can be
made.
Can you suggest why the first Nobel gas compound
prepared contained Xe?
Questions for you to try:
1. Explain why the third ionisation energy of
Magnesium is so much greater than its second.
2. Calculate the energy change for the following
changes.
a)Ca (g) → Ca2+ (g) + 2eb)B2+(g) → B4+ (g) + 2e3. Which of the following equations represents the first
ionisation energy of fluorine?
A. F–(g) → F(g) + e–
B. F–(g) → F2(g) + e–
C. F(g) → F+(g) + e–
D. F2(g) → F+(g) + e–
4. Which line in the table is likely to be
correct for the element francium?
State at 30 °C
First ionisation
energy/kJ mol–1
A
solid
less than 382
B
liquid
less than 382
C
solid
greater than 382
D
liquid
greater than 382
5. As the atomic number of the alkali metals increases
A the first ionisation energy decreases
B the atomic size decreases
C the density decreases
D the melting point increases.
6. Explain why
a) A potassium atom is larger than a sodium atom
b) The Chlorine atom is smaller than a sodium
atom
7. Atoms of different elements are different sizes.
What is the trend in atomic size across the period
from sodium to argon?
8. Which of the following reactions refers to the
third ionisation energy of aluminium?
A Al(s) → Al3+(g) + 3e–
B Al(g) → Al3+(g) + 3e–
C Al2+(g) → Al3+(g) + e–
D Al3+(g) → Al4+(g) + e–
9. Atoms of different elements have different
ionisation energies.
Explain clearly why the first ionisation energy of
potassium is less than the first ionisation energy of
sodium.
Answers
1.Explain why the third ionisation energy of
Magnesium is so much greater than its second.
Removing the third mole of electrons involves breaking into a
stable complete energy level of electrons. This energy level
is closer to the nucleus so more attraction between the
outer electron and nucleus.
2. Calculate the energy change for the following
changes.
a)Ca (g) → Ca2+ (g) + 2e-
596 + 1160 = 1756 kJmol-1
b)B2+(g) → B4+ (g) + 2e3660 + 25000 = 28660 kJmol-1
3. Which of the following equations represents
the first ionisation energy of fluorine?
A F–(g) → F(g) + e–
B F–(g) → F2(g) + e–
C F(g) → F+(g) + e–
D F2(g) → F+(g) + e–
C
4. Which line in the table is likely to be
correct for the element francium?
State at 30 °C
First ionisation
energy/kJ mol–1
A
solid
less than 382
B
liquid
less than 382
C
solid
greater than 382
D
liquid
greater than 382
B
5. As the atomic number of the alkali metals increases
A the first ionisation energy decreases
B the atomic size decreases
C the density decreases
D the melting point increases.
A
6. Explain why
a) A potassium atom is larger than a sodium atom
Potassium has an extra energy level (shell) of
electrons.
b) The Chlorine atom is smaller than a sodium
atom
Both atoms have the same number of energy levels,
but the chlorine has a greater nuclear charge than
sodium. This attracts the outer electrons more
strongly and causes the decrease in atomic radius.
7. Atoms of different elements are different sizes.
What is the trend in atomic size across the period
from sodium to argon?
Decreases or gets smaller
8. Which of the following reactions refers to the
third ionisation energy of aluminium?
A Al(s) → Al3+(g) + 3e–
B Al(g) → Al3+(g) + 3e–
C Al2+(g) → Al3+(g) + e–
D Al3+(g) → Al4+(g) + e–
C
9. Atoms of different elements have different
ionisation energies.
Explain clearly why the first ionisation energy of
potassium is less than the first ionisation energy of
sodium.
Potassium has more electron shells (or outer electron is
further from the nucleus)
The inner electrons (electron shells) shield (screen) the
outer electron from the attraction of the nucleus.
Therefore the outer electron is held less tightly in
potassium
Periodic Table of Data Visual database of
the physical and thermochemical
properties of the chemical elements which
allows the user to plot graphs and tables,
play games and view diagrams.
Periodic Landscapes The Periodic Landscape
images are computer-generated landscape
views and models based on patterns and
relationships within the periodic table. The
models are sculpted to achieve a sense of
general trends or patterns.
Bonding and structure in
the first 20 elements
Overview
Learn how the first twenty
elements can be classified into
groups according to their
bonding and structure.
Chemical Bond
Intra-molecular
(Metallic, ionic and
covalent bonding)
Inter-molecular
(between molecules)
Van der Waals’
http://www.educationscotland
.gov.uk/highersciences/chemi
stry/animations/intermolecula
rforces.asp
This animation describes and explains the key intermolecular
forces of attraction (Van der Waals forces of attraction)
including London dispersion forces, permanent dipole-permanent
dipole attractions and Hydrogen Bonding.
Types of bonding in elements
• Metallic
• Covalent
• Both of these are intra-molecular
forces – i.e. bonds between atoms!
The Metals
Metallic bonding
Groups I, II, III
Metallic elements
Delocalised
electron
+
+
+
+
+
+
+
+
Positive nucleus (core)
Electron shells
The outer shell in metals is not full and so the outer electrons in
metal atoms can move randomly between these partially filled
outer shells. The electrons are delocalised (sometimes called a
‘sea’ or ‘cloud’ of
Electrons) i.e. they are held in common by all the atoms.
Metallic bonding is the strong electrostatic force of attraction,
between the positive charged ions, formed by the loss of the
outer shell electrons of a metal atom and these delocalised
electrons.
The positive metal ions are held together by this
electron “Glue”.
The outer electrons are delocalised and free to move
throughout the lattice.
The greater the number of electrons in the outer
shell the stronger the metallic bond.
So the melting point of Al>Mg>Na
Conduction of Metals
Metallic Character
The strength of the metallic bond depends on
1) The elements tendency to lose electrons (ionise)
2) The packing arrangement of the metal atoms.
3) The size of the atom.
4) The number of valence electrons in the outer most
shell.
5) The number of shells.
Physical properties of metals
A. Metals are malleable and ductile
Metal atoms can ‘slip’ past each other because the metallic
bond is not fixed and it acts in all directions.
Physical properties of metals
B. Conduction of electricity and thermal energy.
Solid and liquid metals conduct heat and electricity.
The delocalised electrons are free to move in the
solid lattice. These mobile electrons can act as
charge carriers in the conduction of electricity or
as energy conductors in the conduction of heat.
Physical properties of metals
C. Change of state
In general, metals have high melting and boiling points
because of the strength of the metallic bond.
When a metal is molten the metallic bonds are still
present.
B.p.’s are much higher as you need to break the metallic
bonds throughout the metal lattice.
Metal boiling point trends
Down a group
Boiling point of alkali metals
1600
Boiling point /oC
1400
1200
1000
800
Series1
600
400
200
0
Lithium
Sodium
Potassium
Metal
Boiling point across a period
Boiling point /oC
3000
2500
2000
1500
Series1
1000
500
0
Potassium
Calcium
Metal
Gallium
The strength of metallic
bonding decreases the
forces of attraction get
weaker.
Across a period
The covalent radius
decreases as the positive
core is increasing in charge,
this has the effect of pulling
the outer closer to the
nucleus the forces of
attraction increase.
The Non-metals
Nobel gases
Group 0
Noble gases
He
Noble gases have full outer electron shells
++
They do not need to combine with other atoms.
They are said to be monatomic.
Group 0 are all gases and
exist as individual atoms.
However, the monatomic gases do form weak inter-atomic
bonds at very low temperatures.
London dispersion forces
Learning intention
Learn
about
these
weak
intermolecular forces of attraction
– they exist between molecules of
all non-metal discrete molecular
elements and the atoms of
monatomic noble gases.
London Dispersion Forces
Monatomic elements
++
Sometimes the electrons can end up on one
side of the atom, i.e. the electron cloud can wobble
++
This means that one side of the atom is more
negative than the other side.
δ+
δ-
δ-
δ+
These charges are given the symbol δ ‘delta’
A temporary dipole is therefore formed.
δ+
A dipole can induce other atoms
to form dipoles, resulting in
dipole –dipole attraction.
δLondons forces
Monatomic elements
London dispersion forces are very weak attractive forces
Noble gases b.p.’s
180
160
166
140
120
121
100
b.p / K
80
87
60
40
20
4
Helium
Neon
Argon
Krypton
Xeon
27
0 K = -293o C 0
B.p.’s increase as the size of the atom increases.
This happens because the London forces increases with
increasing number of electrons. The more electrons the
bigger the temporary dipole, the stronger the London
forces – so higher b.p.
The Non-metals
1. Discrete Covalent molecules
Relating physical properties to
intermolecular forces in
Discrete Covalent molecules
Learning intention
Learn how to explain differences
in physical properties such as
viscosity, melting point and boiling
point in terms of differences in
strength of intermolecular forces.
Covalent molecular elements
Most non-metals exist as discrete covalent
molecules where their atoms are held together by
strong covalent bonds.
e.g. O2,S,CL2,P,N2,F2......
Discrete molecules have a definite formula with a
definite number of atoms bonded together. They are
very small(discrete) molecules.
The discrete molecules are held by weak
intermolecular forces called London Dispersion
Forces. This means they have .......?..... Melting and
boiling points.
Covalent Bond - Revision
Chlorine atom
2,8,7
17+
Chlorine molecule Cl2
2,8,8
17+
17+
diatomic
A covalent bond is formed when a pair of electrons
are shared.
The atoms in a covalent bond is the mutual
attraction of two positive nuclei for a shared pair of
electrons. This makes it very strong.
Group IV
Buckminster fullerene (Bucky Balls) were discovered
in the 1980’s.
C60
C70
Due to the large molecules , fullerenes have stronger
London forces between their molecules, compared to
elements made from smaller molecules.
Fullerenes are a family of carbon molecules made up
of rings with definite formula.
They are discrete covalent molecules
Group V
Strong triple
covalent bond
N
Weak Londons force
N
N
N
Strong intra-molecular bonding and
weak inter-molecular bonding exist
in this diatomic molecule.
N 2 m.p. -210o C
Phosphorus P4
m.p. 44oC
Strong covalent
bonds
Weak Londons
forces
Group VI
Strong double
covalent bond
O
Weak Londons force
O
O
O
Strong intra-molecular bonding and
weak inter-molecular bonding exist
in this diatomic molecule.
O 2 m.p. -218o C
Sulphur S8
Weak Londons
forces
m.p. 113oC
Higher m.p. because
there arestronger
Londons’ forcesbetween
larger molecules(more
electrons).
Group VII
Fluorine F2
Strong covalent
bond
F
Weak London forces
Strong intra-molecular bonding
(Covalent) and weak inter-molecular
bonding (London forces) exist in
this diatomic molecule.
F 2 m.p. -220o C
F
F
F
Chlorine Cl2
Cl
Strong covalent
bond
Weak London forces
Cl
Cl
Cl
Strong intra-molecular bonding (?)
and weak inter-molecular bonding
(?) exist in this diatomic molecule.
Cl 2 m.p. -101oC
Halogens b.p.’s
500
450
457
400
b.p./ K
Fluorine
350
Chlorine
300
332
250
200
238
Bromine
Iodine
150
100
50
85
0
As the size of the halogen atom increases (more electrons),
so does the size of the London forces between the halogen
molecule.
Covalent Network Elements
In the first 20 elements, only Boron, Carbon and
Silicon have covalent network structures.
Carbon Diamond
Diamond forms an infinite 3D
network structure.
Each carbon atom forms 4
covalent bonds to 4 other
carbon atoms.
Very rigid strong structure.
Diamond is one of the hardest
materials known to man.
C sublimes 3642oC
Graphite
Carbon bonded to only 3 other Carbons
So the spare electrons are delocalised
and so free to move. Graphite is a conductor.
London forces between the
layers allows layers to slide over
each other.
Graphite can be used as a lubricant
Silicon
Silicon has the same infinite 3D network
structure as diamond Si mp 1410oC.
Trending in Properties –
Revisited
The link to bonding
Density change across a period
3
2.5
Density
g/cm3
Sodium
Magnesium
Aluminium
Silicon
Phosphorus
Sulphur
Chlorine
Argon
2
1.5
1
0.5
0
Na Mg Al Si P
S Cl Ar
Na to Al the atom size decreases leading to greater packing in metal latt
Si is a covalent network, tightly packed atoms in covalent lattice.
P and S are covalent molecular solids with quite densely packed
molecules.
Cl and is a covalent molecular gas at room temperature.
Ar and is a monomolecular gas at room temperature.
Bonding in the first 20 elements
Covalent molecular gases
Li
These elements occur as diatomic
Covalent networks
(two atom) molecules with strong
Giant network
of atoms
withbetween
strong the atoms
covalent
bonds
covalent bonds
between the atoms.
(intramolecular
bonds) and weak
Londons
Very high melting
andforces
boilingbetween
points. the
molecules (intermolecular bonds).
The weak Londons forces mean low
melting and boiling points.
Be
CC
N
F
B
O
Na
Mg
K
Ca
H
Metallic bonding.
Giant network of positively
charged nuclei surrounded
by delocalised electrons.
Delocalised electrons make
these elements good
conductors.
Al
Si
P
Cl
S
He
Ne
Ar
Covalent molecular solids
Polyatomic (many atom) molecules.
Fullerenes C60 C70
Monatomic elements
P4 and S8. These molecules have many
The
noblelarger
gasesLondon
exist as individual
electrons and this
produces
(monatomic)
atoms. There are only weak
forces than the diatomic
molecules.
Londons forces between the atoms.
The stronger London forces (temporary
Very
energy
is needed to break
dipoles) gives these
twolittle
elements
higher
these
forces and so the noble gases
melting and boiling
points.
have very low melting and boiling points.
These two elements are solids at room
temperature.
Diagram shows part of the
covalent
network
of carbon
Uneven
distribution
of the electrons in
atoms
in
diamond.
the electron cloud create temporary
dipoles
(d atom
+ and is
d-)covalently
which result in a weak
Each
carbon
attraction
between
atoms which come
bonded
to 4 other
carbon
close
to
each
other.
These weak
atoms.
attractions are called Londons forces.
d+
dd+
d-
H
He
Li
Be
B
CC
N
O
F
Ne
Na
Mg
Al
Si
P
S
S
Cl
Ar
K
Ca
Metallic bonding with a network of
positively charged nuclei
surrounded by a ‘sea’ of
delocalised electrons.
Diagram
Strong
showscovalent
S8
bonds
molecules
between
in sulphur
the atoms inside
-thelondons
- forces
-molecules.
with the
diatomic
shown
by
the
dotted
+
+
+
+
+
+
+
lines.
+
+
+
+
+
+
+
These large
Weakmolecules
van der Waal’s
have +stronger
+
+ Londons
+
+ molecules
+
+
forces
between
forces than
the
which come close to each
diatomic
molecules.
-other.
- - - -
http://www.educationscotland.gov.uk/highersciences/c
hemistry/animations/bondingstructure.asp
This interactive animation provides a visual
representation of the bonding and structure of the
first twenty elements in the periodic table, taking into
account both the intra- and inter-molecular forces
involved.
Questions on elements – bonding and
structure
1. Explain why the covalent network elements have
high melting and boiling points.
2. Explain why the discrete molecular and monatomic
elements have low melting and boiling points.
3. Does diamond conduct electricity? Explain.
4. Does graphite conduct electricity? Explain.
5. How does the hardness of diamond compare with
graphite? Explain.
6. Give a use for both diamond and graphite.
7. Complete the following table:
Questions on elements – bonding and structure
7. Complete the following table:
Type of bonding and
structure
Metallic solids
Properties
……………. of electricity
Covalent network solids
……….. …. melting points
……………. of electricity
exception ……………….
Covalent molecular
solids
………….. melting points
…………… of electricity
Covalent molecular
(diatomic) gases
and monatomic gases
…………… boiling points
b) Enthalpies of combustion
Learning intention
Learn the definition of enthalpy of
combustion, which can be directly
measured using a calorimeter.
National 5 = Energy of Combustion
Higher = Enthalpy(ΔH) of combustion
• Definition: The enthalpy of combustion is the heat
energy given out when 1 mole of fuel burns
completely in oxygen.
• The enthalpy of combustion of methane can be
represented by the equation
• CH4(g) + O2 (g)
CO2(g) + H2 O(l)
Enthalpy of combustion
The heat energy released when alcohols burn can be measured
The enthalpy of combustion of a substance is the amount of energy
given out when one mole of a substance burns in excess oxygen.
Calculating Enthalpy of CombustionSpecific
heat capacity
Calculating the energy change during a chemical reaction in water.
Calculation
(a)The heat energy gained by the water (Eh) can be calculated
using the formula:
H = c. m.  T
c
m
T
=specific heat capacity
=mass in Kg
=temperature change
The mass of water can be calculated by using the fact that 1 ml = 1 g.
The value for c is usually taken as 4.18 kJ kg –1 oC-1
Measuring the enthalpy of combustion of alcohols
A Procedure?
Weigh a filled alcohol burner
Measure 100 cm3 water into a copper calorimeter
Take temperature of the water
Light the burner and use it to heat the water for 3 minutes.
Stir the water and take the highest temperature reached
Reweigh the burner and remaining fuel
Calculate the energy for mass burned
Calculate the Enthalpy for 1 mole of the alcohol
Enthalpy of combustion
Procedure
1. Weigh the spirit burner (already containing ethanol) with its cap on and
record its mass. (The cap should be kept on to cut down the loss of
ethanol through evaporation)
2. Using the measuring cylinder, measure out 100 cm3 of water into the
copper can.
3. Set up the apparatus as directed by your teacher/lecturer.
4. Measure and record the temperature of the water.
5.Remove the cap from the spirit burner and immediately light the burner.
6.Slowly and continuously stir the water with the thermometer. After 3
minutes, recap the spirit burner and measure and record the maximum
temperature of the water.
7. Reweigh the spirit burner and record its mass.
CALCULATION
Suppose 0.25 g of ethanol had been burned and the temperature of 100cm3 water had risen
by 12.5 °C.
Part 1
The heat energy gained by the water (Eh) is calculated using the formula:
Eh = c m T
Eh = 4.18 x 0.10 x 12.5
= 5.225 kJ
We assume that the heat energy released by the burning ethanol is gained only by the water.
The heat energy released on burning 0.25 g of ethanol = 5.225 kJ
However Enthalpy is for 1 mole of a substance so……….
Part 2
Work out GFM (1 mole in grams) of alcohol…….
Ethanol: CH3CH2OH
Mass of 1 mole = 2(12) + 6(1) + 16 = 46 g
We can now calculate the heat energy released on burning 1 mole of ethanol. (The Enthalpy of
Ethanol)
0.25g
5.225kJ x 5.225
46g
46
=0.25
961 kJ
The enthalpy of combustion of ethanol = - 961 kJ mol-1
(A negative sign is used because combustion is an exothermic reaction)
Measuring the enthalpy of combustion of alcohols
The heat energy gained by the water (Eh) is calculated using the formula:
Eh
= c m ∆T
Eh
=
x
x
=
kJ
Measuring the enthalpy of combustion of alcohols
We assume that the heat energy released by the burning alcohol is
gained only by the water.
The heat energy released on burning ……….. g of ……………anol
So one mole .............. g of ................anol
………….. kJ
....................kJ
The enthalpy of combustion of …………anol = -………….. kJ mol-1
(A negative sign is used because combustion is an exothermic reaction)
Sources of inaccuracy
• Heat loss to surroundings
• Incomplete combustion
• Possible loss of fuel by evaporation
from wick
Calorimetry
To eliminate these inaccuracies a bomb
calorimeter is used
The burning fuel (or food) is supplied with
oxygen to encourage complete
combustion
The combustion chamber is entirely
surrounded so there is no heat loss to
the surroundings
Commercial ‘bomb’ calorimeters
Databooklet Values
The calorimeter is heated electrically.
Energy required to heat the entire
apparatus by 1 0C is calculated.
Enthalpy of combustion
Worked example 1.
0.19 g of methanol, CH3OH, is burned and the heat energy given out
increased the temperature of 100g of water from 22oC to 32oC.
Calculate the enthalpy of combustion of methanol.
–704 kJ mol-1
Enthalpy of combustion
Worked example 1.
0.19 g of methanol, CH3OH, is burned and the heat energy given out
increased the temperature of 100g of water from 22oC to 32oC.
Calculate the enthalpy of combustion of methanol.
–704 kJ mol-1
Worked example 2.
0.22g of propane was used to heat 200cm3 of water at 20oC. Use the
enthalpy of combustion of propane in the data book to calculate the
final temperature of the water.
33.3oC
Calculations for you to try.
1.
0.25g of ethanol, C2H5OH, was burned and the heat given out raised
the temperature of 500 cm3 of water from 20.1oC to 23.4oC.
Calculate the enthalpy of combustion of ethanol
2. 0.01 moles of methane was burned and the energy given out raised
the temperature of 200cm3 of water from 18oC to 28.6oC. Calculate
the enthalpy of combustion of methane.
3. 0.1g of methanol, CH3OH, was burned and the heat given out used to
raise the temperature of 100 cm3 of water at 21oC.
Use the enthalpy of combustion of methanol in the data booklet to
calculate the final temperature of the water.
4. 0.2g of methane, CH4, was burned and the heat given out used to
raise the temperature of 250 cm3 of water
Use the enthalpy of combustion of methane in the data booklet
to calculate the temperature rise of the water.
1.
0.25g of ethanol, C2H5OH, was burned and the heat given out raised
the temperature of 500 cm3 of water from 20.1oC to 23.4oC.
Calculate the enthalpy of combustion of ethanol -1269 kJ mol-1
2. 0.01 moles of methane was burned and the energy given out raised
the temperature of 200cm3 of water from 18oC to 28.6oC. Calculate the
enthalpy of combustion of methane.
-1
-886.2 kJ mol
3. 0.1g of methanol, CH3OH, was burned and the heat given out used to
raise the temperature of 100 cm3 of water at 21oC.
Use the enthalpy of combustion of methanol in the data booklet to
calculate the final temperature of the water. 26.4oC
4. 0.2g of methane, CH4, was burned and the heat given out used to
raise the temperature of 250 cm3 of water
Use the enthalpy of combustion of methane in the data booklet
to calculate the temperature rise of the water. 10.66oC
Tutorial Question – Self Check
Page 10 of Tutorial booklet;
1. ΔH of butanol = -2664 kjmol -1
2. ΔH of ethanol = -1345 kjmol -1
3. ΔH of sulphur = -295 kjmol -1
4. ΔH of methane = - 877 kjmol -1
Now check your Combustion Chemical
Equations!
b) Calculation of the mass of
products: Balanced Equation
Calculations
Learning intention
Learn how the theoretical mass or
volume of product can be
calculated from the balanced
reaction equation.
From previous studies
You should be able to
• Write Formulae
• Calculate percentage composition
• Calculate the number of moles in a given mass
• Calculate the number of moles of solute
dissolved in a solution
• Use balanced equations to work out unknown
masses
Reacting Masses
Magnesium + oxygen….how much magnesium
oxide is made?
1.
Accurately weigh a crucible
2. Add approx 1.2g of Mg ribbon and reweigh
3. Place the crucible and lid in a silica triangle.
4. Heat gently at first then more strongly. Lift the lid with
tongs from time to time to admit more oxygen, but not enough
to let out the magnesium oxide.
Reacting masses
1. When the reaction is complete, the magnesium
will not glow more brightly when the lid is raised
2. Allow the crucible to cool
3. Reweigh the crucible
Calculations
Using the balanced equation calculate the mass of MgO you
would expect to be formed?
2Mg + O2
→
2MgO
Reacting masses - Questions
1. What mass of zinc sulphate will be produced on adding
6.0g zinc to excess sulphuric acid?
14.9g
2. What mass of sodium carbonate will react completely with
100cm3 of nitric acid concentration 1 mol l-1?
5.3g
d) Excess
Learning intention
Learn how to calculate how much
of a particular reactant is in
excess
from
the
balanced
equation.
Excess
You can use the relative numbers of moles of substances, as shown in
balanced equations, to calculate the amounts of reactants needed or
the amounts of products produced.
A limiting reactant is the substance that is fully used up and
thereby limits the possible extent of the reaction. Other reactants
are said to be in excess.
Calculations involving excess
A limiting reactant is the substance that is fully used up and
thereby limits the reaction going further.
Other reactants are said to be in excess.
As soon as one of the reactant in a chemical reaction is used up the
reaction stops. Any other reactant which is left over is said to be ‘in
excess’. The reactant which is used up fully (limiting reactant)
determines the mass of product formed – SO IS MOST IMPROTANT in
terms of how much product is formed!
You can use the relative numbers of moles of substances (mole
ratios), as shown in balanced equations, to calculate the amounts of
reactants needed or the amounts of products produced.
Calculations involving excess
Starter
a) Which reactant is the limiting reactant when 10g of calcium
carbonate reacts with 100cm3 of 1 mol l-1 hydrochloric acid?
Equation: CaCO3 + 2HCl
→
CaCl2
+
H2O + CO2
b) Now calculate the mass of carbon dioxide formed.
CaCO3 is in excess.
CO2 = O.22g
Calculations involving excess
Worked example 2. 1.2g of magnesium was added to 100cm3 2 mol
l-1 hydrochloric acid. Calculate the reagent in excess and therefore
the limiting reactant?
Acid in excess.
Graphs and Rates of Reaction
e.g. Zn + 2HCl  ZnCl2 + H2
Zn in excess
2mol l-1 HCl 20oC
2mol l-1 HCl 40oC
Faster, but same amount of
gas produced
Vol
H2
cm 3
1 mol l-1 HCl 20oC
Half the gas produced
time /s
HCl is limiting reagent
Calculations for you to try.
1.
The graph below was obtained when 1.0g of powdered zinc was
added to excess hydrochloric acid 1.0 mol l-1, copy the graph and
sketch a line to show what you would expect if the reaction was
repeated using
a) 2.0 mol l-1 HCl and 1.0g Zn
b) 1.0 mol l-1 HCl and 0.75g Zn
Vol
H2
cm 3
time /s
2. a) Calculate the limiting reactant when 3.27g of zinc is reacted with
100cm3 of 2.0 mol l-1 hydrochloric acid.
b) What mass of hydrogen gas will be produced?
0.1 g
Calculations involving excess
Examples. For each of the following reactions calculate which
reagent is in excess?
a) 4.86g magnesium added to 250cm3 2 mol l-1 hydrochloric acid
HCl
b) 2.7g aluminium added to 200cm3 1 mol l-1 hydrochloric acid
Al
c) 2.43g magnesium added to 200cm3 1 mol l-1 sulphuric acid
H2SO4
d) 3.27g zinc added to 100cm3 0.2 mol l-1 hydrochloric acid.
Zn
Calculations involving excess and molar gas volume
An experiment was carried out to measure the concentration
of hypochlotite ions (ClO-) in a sample of bleach. In this
experiment the bleach sample is reacted with excess hydrogen
peroxide.
H2O2(aq) + ClO-(aq) H2O(l) + Cl-(aq) + O2(g)
By measuring the volume of oxygen given off, the
concentration of the bleach can be calculated.
80cm3 of oxygen was produced from 5.0cm3 of bleach.
Calculate the concentration of the hypochlorite ions in the
bleach (Take the molar gas volume to be 24 Litres mol-1).
Calculations for you to try.
1. What mass of calcium oxide is formed when 0.4 g of calcium reacts
with 0.05 mole of oxygen?
2Ca +
O2

2CaO
0.56g
2. What mass of hydrogen is formed when 3.27g of zinc is reacted with
25cm3 of 2 mol l-1 hydrochloric acid?
Zn
+ 2HCl

ZnCl2
+ H2
0.05 g
Bonding in compounds
Overview
Learn how the elements can
form bonds in compounds.
Van der Waals forces
Learning intention
An introduction to the variety
of intermolecular forces which
exist between molecules.
Relating physical properties to
intermolecular forces
Learning intention
Learn how to explain differences in
physical
properties
such
as
viscosity, melting point and boiling
point in terms of differences in
strength of intermolecular forces.
The Chemical Bond
Types of Chemical Bonding in
Compounds
Intramolecular
(Within – i.e. between
atoms)
Intermolecular
(between molecules)
Van der Waals Forces
Note only Covalent molecular compounds
Ionic (compound only)
Covalent bond(element or
compound)
New ! Polar covalent bond
(compound only)
contain these.
1. Hydrogen bonding
2. Permanent Dipole-Permanent
Dipole interactions
3. London dispersion forces
Bond Strengths
Bond Type
Strength (kJ mol –1)
Metallic
80 to 600
Ionic
100 to 500
Covalent
100 to 500
Hydrogen
40
Dipole-Dipole
30
Londons forces
1 to 20
Ionic Compounds
Ionic Compounds
Ions
- metals lose electrons and form positive ions
- non-metals gain electrons to form negative ions
- electrons are transferred from metals to non-metals
transfer
-
+
Na
+
Cl
Na atom + Cl atom
(2.8.1)
(2.8.7)
Na
Cl
Na+ ion + Cl- ion
(2.8)
(2.8.8)
Covalent Molecular
Compounds
Covalent Bonding
Sharing electrons
•
takes place between non-metal and non-metal
•
shared electrons count as part of the outer shell
of both Atoms
•
shared electrons attract the nuclei of both atoms
•
this attraction is called the covalent bond
Hydrogen chloride
H
H
Cl
(linear)
Cl
HCl
Ammonia
H
H
H
N
H
H
(pyrimidal) NH3
H
N
Water
O
H
H
(bent) H2O
H
O
H
Draw electron dot cross diagrams for the following
molecules and structural formula
1. SCl2
2. CO2
3. CH4
X X
S
X X
H
Cl-S-Cl
O=C=O
H C H
H
Structure of Ionic
Compounds
Ionic Compounds
The positive and negative ions are attracted
(electrostatic bond )to each other.
Ionic bond (electrostatic attraction)
Na+
Cl-
A giant lattice structure is formed.
Each Na+ ion is surrounded by 6 Cl- ions.
While each Cl- is surrounded by 6 Na+ ions.
Ionic bonding is the electrostatic force of
attractionbetween positively and negatively
charged ions.
This ionic network compound has many ionic
bonds so ionic compounds have high m.p.s
Ionic Compounds
A giant lattice structure is formed when each Na+ ion is
surrounded by 6 Cl- ions and each Cl- ion is surrounded by
6 Na+ ions.
Sodium Chloride
The formula of sodium chloride
is NaCl, showing that the ratio
of Na+ to Cl- ions is 1 to 1.
The m.p. of NaCl is 801 0C
The size of the ions will effect the strength of the ionic bond
and how the ions pack together. e.g. NaF m.p. 1000oC, NaI 660oC
Molecular Ions, e.g. SO4
Oxygen
A single covalent bond.
Sulphur
O
2 additional electrons
e.g. Copper can donate
the extra 2 electrons
needed.
Cu
Cu
2+
+ 2e
O
S
O
O
Copper sulphate contains the Cu2+ and the SO42- ions. There is,
therefore, covalent bonding and ionic bonding in copper sulphate
A solution of copper sulphate can conduct electricity.
Molten ionic compounds can also conduct electricity.
Bond Strengths
Bond Type
Strength (kJ mol –1)
Metallic
80 to 600
Ionic
100 to 500
Covalent
100 to 500
Hydrogen
40
Dipole-Dipole
30
London’s forces
1 to 20
Discrete Covalent Molecular
Compounds
Covalent Molecular Compounds
Discrete molecules are formed when two or more atoms share
electrons.
The atoms are non-metal elements. An example is methane.
Methane: CH4
H
H
H
C
H
H
H
C
H
H
Methane has strong intra-molecular and weak inter-molecular.
It’s b.p. is -183oC
Covalent Molecular Compounds
Non- metals elements can form double and triple covalent bonds.
ethane C2H6
H
H
H
C
C
H
H
H
ethene C2H4
H
H
H
C
C
H
H
H
H
H
C
C
H
H
H
H
C
C
H
H
H C
Double covalent bond
Covalent molecular compounds have low m.p.’s because the weak
forces holding the molecules together require only small
amounts of thermal energy to break them.
Bond Strengths
Bond Type
Strength (kJ mol –1)
Metallic
80 to 600
Ionic
100 to 500
Covalent
100 to 500
Hydrogen
40
Dipole-Dipole
30
London’s Forces
1 to 20
Covalent Molecular Compounds
Properties
Low m.p.’s and b.p.’s., this increases with size of the molecule
and the increasing number of atoms in the molecule.
m.p.’s of the carbon halides
171
Temp
/ oC
90
-23
-183
CF4
CCl4
CBr4
CI4
m.p.’s increase because the strength of the London dispersion
forces increase with the increasing size of the molecule. So more
Energy is needed to separate molecules.
Covalent Network
Compounds
Silicon Dioxide SiO2
Silicon and oxygen make up nearly 75% of the Earth’s crust.
They are therefore the most common elements in the
Earth’s crust.
They combine together to make a covalent network compound
called silicon dioxide.
This is usually found in the form of sand or quartz.
Each Si atom is bonded to 4 O atoms, and each O atom is
bonded to 2 Si atoms. Hence the chemical formula, SiO2 .
Silicon dioxide (silica) also has a
high m.p. (1713 oC) and like SiC,
it is very hard and used as an
abrasive.
It is relatively un-reactive.
New Higher Chemistry E Allan J Harris
Silicon Carbide SiC
Silicon, like carbon, can form giant covalent networks.
Silicon carbide exist in a similar structure to diamond.
Tetrahedral
shape
C
Covalent
Bond
C
Si
C
C
The 4 carbon atoms are
available to bond with
another 4 silicon atoms.
This results in a COVALENT NETWORK COMPOUND
Silicon Carbide SiC
Silicon carbide (carborundum) has a
chemical formula is SiC. As this compound is
linked by strong covalent bonding, it has a
high m.p. (2730oC).
It is a hard substance as it is very difficult
to break the covalent lattice.
SiC is used as an abrasive for smoothing
very hard materials.
Each Si is bonded to 4 C’s and
each C is bonded to 4 Si’s.
Hence the chemical formula, SiC
3.32 Periodic trends in
electronegativity
Learning intention
• Learn
the
definition
of
electronegativity
• how to explain periodic trends in
terms of nuclear charge, covalent
radius and the screening effect of
inner electrons.
Electronegativity
Electronegativity is a numerical measure of the relative
ability of an atom in a molecule to attract the bonding
electrons towards itself.
Electronegativity
Electronegativity is a measure of an atom’s
attraction for the shared pair of electrons
in a bond.
e
C
e
H
Which atom would have a greater
attraction for the electrons in this
bond and why?
Linus Pauling
Linus
American
TodayPauling,
we stillanmeasure
chemist
(and winner of
of two
electronegativities
Nobel
prizes!)
up with the
elements
usingcame
the Pauling
concept
scale. of electronegativity in
1932 to help explain the nature
of chemical bonds.
Since fluorine is the most
electronegative element (has
the greatest attraction for the
Values for electronegativity
bonding electrons) he assigned
can be found on page 11 of
it a value and compared all
the data book
other elements to fluorine.
Electronegativities
Looking across a row or down a group of the
periodic table we can see a trend in values.
We can explain these trends by applying the
same reasoning used for ionisation energies.
Looking across a period
Increasing Electronegativity
Li Be
1. 1.5
0
B C N O F
2. 2. 3. 3. 4.
0 5 0 5 0
What are the
electronegativities of
these elements?
Across a period electronegativity increases
The charge in the nucleus increases across a period.
Greater number of protons = Greater attraction for bonding
electrons
Decreasing
Electronegativity
Looking down a group
F 4.0
Cl 3.0
What are the
electronegativities of
these halogens?
Br 2.8
I 2.6
Down a group electronegativity decreases
Atoms have a bigger radius (more electron shells)
The positive charge of the nucleus is further away from
the bonding electrons and is shielded by the extra
electron shells.
Trends in electronegativity
Electronegativity increases across a period.
Electronegativity decreases down a group
Going across the period, the nuclear charge increases.
This pulls the electron shells closer to the nucleus.
As a results, the electronegativity increases.
Going down the group, the nuclear charge increases but
the number of electron shells also increases.
As a result of ‘shielding’ and an increase distance the outer
shell is from the nucleus, electronegativity decreases.
Chemical bonds: types of bonds Explores how
different types of bonds are formed due to
variations in the electronegativity of the bonded
atoms. The distortion of the orbitals and the
polarity of the bond is also displayed.
Linus Pauling (1901-1994) An account of the life
and work of the Nobel Prize-winning chemist,
Linus Pauling.
Periodic Table of Data Visual database of the
physical and thermochemical properties of the
chemical elements which allows the user to plot
graphs and tables, play games and view diagrams.
The Bonding Continuum
Predicting Non-Polar, Polar
and Ionic Bonds
Learning intention
Learn how differing differences
in electronegativity between
bonding atoms lead to the
formation of non-polar covalent
bonds, polar covalent bonds and
even ionic bonds.
Bonding continuum
Learning intention
Learn
about
the
bonding
continuum
which
stretches
between pure covalent and ionic
bonds, in terms of differences
of electronegativity between
bonding atoms.
Covalent Bonding
A covalent bond is a shared pair of electrons
Both nuclei try to pull the electrons towards
electrostatically attracted to the positive nuclei
themselves
of two atoms.
+
-
+
The
achieve a stable
This
is atoms
like a tug-of-war
whereouter
bothelectron
sides are
arrangement
gas arrangement)
by
pulling(aonnoble
the same
object.
sharing
It creates a strong
bondelectrons.
between the two atoms.
Covalent Bonding
Picture a tug-of-war:
If both teams pull with the same force the midpoint of the rope will not move.
Pure/Non-Polar Covalent Bonds
No/Very small difference on Electronegativity
H
e
e
H
This even sharing of the rope can be compared to a
pure covalent bond, where the bonding pair of
electrons are held at the mid-point between the
nuclei of the bonding atoms. All non-metal
elements have Pure/Non-polar covalent bonds.
Polar Covalent Bonds
Non-polar covalent bond –
electrons shared equally
between atoms (same
electronegativity)
Polar covalent bond –
electrons shared unequally
between atoms (atom B is
more electronegative)
Polar Covalent Covalent Bonds
Bigger difference in electronegativity
What if it was an uneven tug-of-war?
The team on the right are far stronger, so will pull
the rope harder and the mid-point of the rope will
move to the right.
Polar Covalent Bond
A polar covalent bond is a bond formed when the
shared pair of electrons in a covalent bond are not
shared equally.
This is due to different elements having different
electronegativities.
Polar Covalent Bond
δ-
e.g. Hydrogen Iodide
δ+
H
e
e
I
If hydrogen iodide contained a pure covalent
bond,makes
the electrons
would negative
be shared
equally
as
This
iodine slightly
and
hydrogen
shown
above.
slightly
positive. This is known as a dipole.
However, iodine has a higher electronegativity and
pulls the bonding electrons towards itself
(winning the tug-of-war)
Polar Covalent Bond
In general, the electrons in a covalent bond are not
equally shared.
e.g.
C
δ+
2.5
Cl
δ-
3.0
Electronegativities
δ- indicates where the bonding electrons are most
likely to be found.
Polar Covalent Bond
Consider the polarities of the following bonds:
Difference
Electronegativities
Bond
C
Cl
2.5
3.0
0.5
P
H
2.2
2.2
0
O
H
3.5
2.2
1.3
P
H
δ+
C
δ-
Cl
δ-
O
δ+
H
Increasing Polarity
Complete a similar table for C-N, C-O and P-F bonds.
Polar Covalent Bonds
In the covalent bond between
fluorine and hydrogen. The
bonding electrons are not
shared equally between the two
atoms.
Fluorine
Hydrogen
The fluorine nucleus has more
protons and has a stronger pull
on the electrons than the
hydrogen nucleus..
d-
F
4.0
d+
H
2.2
Thus the fluorine atom has
a greater share of the
bonding electrons and
acquires a slight negative
charge.
The hydrogen atom is then made slightly positive.
The bond is a polar covalent bond and we use the
symbols d+ and d- to show this.
The dipole produced is permanent.
Fluorine is the most electronegative element. It is small
atom compared to others and its nucleus is massive for
its atomic size.
Some other polar covalent bonds are O-H and N-H.
They have large differences in Electronegativty.
Bonding Continuum
“Covalent compounds are formed by
non-metals only”
IS NOT AN ABSOLUTE LAW!
Some compounds break this rule….
The greater the difference in electronegativity the
greater the polarity between two bonding atoms and the
more ionic in character.
Electronegativity Difference and Bond Type:
Difference
Bond
Example
0.0-0.5
Non-Polar/(pure)Covalent Bond H-H
0.0
0.5-1.5
Polar Covalent Bond
H-Cl
H20
0.9
0.7
> 1.5
Ionic
NaCl
2.1
*Note Exceptions: Very Polar bonds
1.5-2.0 Very Polar Covalent (almost ionic) H-F
1.9
Making Tin(IV)iodide
• Gently heat the tin
and iodine in a small
conical flask containing
10cm3 of tolulene on a
hot plate.
• Collect the yellow
precipitate by
filtration using
Büchner filtration
Making Tin(IV)iodide
• Determine the
melting point of
the solid collected.
Making Tin(IV)iodide
Melting point of tin(IV)iodide is
143oC.
Tin electronegativity of 1.8
Iodine has electronegativity
of 2.6
Molecule contains polar covalent
bonds, but the symmetry cancels out
the dipoles, therefore only weak
London dispersion forces so low
melting an boiling point.
Titanium (IV) chloride
TiCl4 is a dense, colourless liquid.
It is one of the rare transition metal halides that is
a liquid at room temperature, This property reflects
the fact that TiCl4 is ………….; that is, each TiCl4 ……….
is relatively …………… associated with its neighbours.
Most metal chlorides are ionic. The attraction
between the individual TiCl4 molecules is weak,
primarily ……………….……….., and ……………. these weak Van
der Waals (intermolecular) interactions result in low
melting and boiling points.
TiCl4 is soluble in toluene and dichloromethane, as
are other non-polar species.
TiCl4
• Used in smoke grenades and for
smoke screens
Non-Polar/Polar Bonds vs
Non-Polar/Polar Molecules
• We have just learned about polar and nonpolar bonds and how to identify them.
• But just because a molecule has polar bonds,
doesn’t mean it is overall a polar molecule.
• Checks have top be done!
Symmetry CCl4
Cl d Symmetrical molecule
d+
Cl
Cl
d
d
Cl
Tetrachloromethane
has a symmetrical
d
arrangement of polar
4 polar covalent
C-Cl
bonds and the polarity
bonds in CCl4 tetrahedral
cancels out over the
shape
molecule.
NON-POLAR molecule
Symmetry CO2
Symmetrical molecule
d-
O
d+
O d-
2 polar covalent C=O
bonds in CO2 linear shape
NON-POLAR molecule
Carbon dioxide has a
symmetrical
arrangement of polar
bonds and the polarity
cancels out over the
molecule.
Symmetry CHCl3
Check 1:
It has polar bonds
Check 2:
It is not symmetrical
H
Permanent dipole
Asymmetrical molecule
d+
Cl
d-
Cl
Cl
d-
d-
3 polar covalent C–Cl bonds and
1 polar covalent C-H bond in CHCl3
It is overall a POLAR molecule
London Dispersion Forces
These Only exist in Non-Polar Molecules.
Remember; London Dispersion forces are weak
forces of attraction between molecules.
•They only exist between non-polar molecules.
•All non-metal discrete covalent molecular
molecules are non-polar
•We now know that many covalent molecular
compounds are non-polar
•Non Polar molecules have low m.p. And b.p due to
the weak London forces between the molecules.
A dipole can induce other atoms
to form dipoles, resulting in temporary
dipole –dipole attraction.
Londons forces
Permanent dipole-permanent
dipole interactions
Learning intention
Learn about this additional
intermolecular
force
of
attraction which exists between
polar molecules.
Permenant Dipole-Dipole Attractions – Only in
Polar Molecules
The differing electronegativities of different atoms in a
molecule and the spatial arrangement of polar covalent bonds
can cause a molecule to form a permanent dipole –
i.e. It is a polar molecule..
H
Permanent dipole
Asymmetrical molecule
d+
Cl
d-
Cl
Cl
d-
d-
3 polar covalent C–Cl bonds and
1 polar covalent C-H bond in CHCl3
POLAR molecule
Permanent Dipole-Dipole
interactions
Molecules with permanent
dipoles attract each other.
The attraction is stronger than Londons forces
Note: Hydrogen bonding is a particular example of
dipole-dipole attractions –we will see this later.
Comparing Properties of Polar and Non Polar
Molecules
Both propanone and butane have the same formula mass of 58
however, butane boils at – 1 oC while propanone boils at 56oC
Propanone is a polar molecule as it has a permanent dipole, so has
polar-polar attraction as well as London’s forces between molecules.
H
H
C
H
d-
O
H
C
C
d+
H
H
b.p. 56 o C
Butane has no permanent dipoles, so only London’s forces
between molecules. So has a lower boiling point.
H
H C
H
H
C
H
H
C
H
H
C H
H
b.p. -1 o C
Hydrogen Bonding
Learning intention
Learn about this strong type of
intermolecular forces which
exists
between
molecules
containing N-H, O-H or F-H
bonds.
Relating physical properties to
intermolecular forces
Learning intention
Learn how to explain differences in
physical properties such as viscosity,
melting point and boiling point in
terms of differences in strength of
intermolecular forces.
Intermolecular - Hydrogen Bonding
Consider the compounds formed between
elements in group 4 of the Periodic table and
hydrogen
The group 4 hydrides are CH4, SiH4, GeH4,
SnH4
They are all covalent molecular so have low
melting points and boiling points.
Boiling Point (K)
Group 4
250
200
150
100
50
0
Group 4
CH4
SiH4
GeH4
SnH4
The boiling point increases as you go
down the group.
As you go down the group the central atom gets
bigger.
There are more electrons so a greater chance of
an uneven distribution of electrons within the
atom.
The London’s forces between the molecules gets
stronger as you go down the group.
More energy is needed to separate the
molecules from each other.
Intermolecular – Hydrogen Bonding
A similar pattern would be expected in the other
covalent molecular hydrides
The group 5 hydrides NH3, PH3, AsH3 and SbH3
The group 6 hydrides H2O, H2S, H2Se and H2Te
The group 7 hydrides HF, HCl, HBr and HI
Boiling Point (K)
Group 5
300
200
Group 5
100
0
NH3
PH3
AsH3
SbH3
NH3, has a higher boiling point than expected.
Boiling Point (K)
Group 6
400
300
200
Group 6
100
0
H2O
H2S
H2Se
H2Te
H2O has a higher boiling point than expected.
Boiling Point (K)
Group 7
400
300
200
Group 7
100
0
HF
HCl
HBr
HI
HF has a higher boiling point than expected.
Intermolecular - Hydrogen Bonding
Boiling Points of Hydrides
Boiling Point (K)
400 H O
2
300
HF
Group 4
Group 5
Group 6
Group 7
NH3
200
100
0
Series Number
It is more difficult to separate NH3, H2O and
HF molecules from each other than expected.
Intermolecular - Hydrogen Bonding
These compounds all have H atoms directly bonded
to very electronegative atoms. But only if
electronegativity beween H and other atom is big
enough, to produce a large dipole, will Hydrogen
bonding be possible.
In HF the H-F bond is polar covalent.
The F has a much higher electronegativity than H.
The pair of shared electrons in the covalent bond
spend more time closer to the fluorine than the
hydrogen.
The H-F bond is very polar.
Hδ+ - Fδ-
Intermolecular - Hydrogen Bonding
The HF molecules can attract each other
Hδ+ - Fδ-
Hδ+ - Fδ-
Hδ+ - Fδ-
This is called hydrogen bonding.
Hydrogen bonding is weak but is stronger
than very weak London forces.
Intermolecular - Hydrogen Bonding
NH3 has H atoms directly bonded to
very electronegative N atoms.
Nd-
Hd+
Hd+ Hd+ Hd+
Nd-
Nd-
Hd+
Hd+ Hd+
Hd+
Hd+
There are
Hydrogen bonds
as well as London
forces between
the ammonia
molecules.
Intermolecular - Hydrogen Bonding
H2O has H atoms directly bonded to
very electronegative O atoms.
Od-
Hd+
Hd+
Od-
Hd+
Od-
Hd+
Hd+
Hd+
There are
Hydrogen
bonds as well
as London
forces
between the
water
molecules.
Proteins consist of long chain atoms containing
polar C=O and H-N bonds.
Hydrogen bonds help give enzymes their shape.
Water
d-
O
H
H
d+
d+
Oxygen has 2 lone pairs of electrons which
can form a hydrogen bonds with two
hydrogen atoms.
Each water molecule, in theory, could be
surrounded by 4 hydrogen bonds.
Hydrogen bonding in water
Density of water
Water
Water has its greatest density at a temperature of 4oC.
When, as water cools further, the molecules start to
move further apart, due to the hydrogen bonding, until a
more open structure is formed at its freezing point. So
ice floats!!
New Higher Chemistry E Allan J Harris
Hydrogen bonding in ice
Hydrogen bonding
in solid water gives
rise to an open
structure. This is
why ice is less
dense than liquid
water.
Hydrogen bonding is also responsible
for holding the two strands of nucleic
acids together in DNA
Viscosity
Density of water
Viscosity is
related to
the molecular
mass and the
number of –
OH present.
Hydrogen
bonding
between the
molecules will
increase its
viscosity.
New Higher Chemistry E Allan J Harris
Water
Surface tension
Water has a high surface tension. The molecules on
the surface have in effect, hydrogen bonds. This
has the effect of pulling the surface molecules
closer together.
Bond Strengths
Bond Type
Strength (kJ mol –1)
Metallic
80 to 600
Ionic
100 to 500
Covalent
100 to 500
Hydrogen
40
Dipole-Dipole
30
London’s forces
1 to 20
Behaviour in electrical fields
New Higher Chemistry E Allan J Harris
Nappies
• Cloth nappies cost between £100-£400 as opposed to
disposable at £800-£1,200 for the 2.5 years of normal
nappy use.
• 3 billion nappies are thrown away in the UK each year
with 90% going to landfill. They can take up to 500
years to decompose.
• Disposables make up 4% of total household waste and
up to 50% of that of families with one baby
• Disposable nappies use up to 5 times more energy to
produce than cotton ones – that's including the washing
process .
• Seven million trees are felled every year in Canada and
Scandinavia to supply the pulp for disposables sold in
the UK.
Sodium polyacrylate is a polymer with a molecular weight of
over one million!
sodium carboxylate
Chemical Background
Groups called sodium carboxylate are attached along the
backbone.
Sodium poly(acrylate) absorbs 500 times its own mass of
water.
water
+
+
Na+
-
Sodium poly(acrylate) absorbs 500 times its own mass of
water.
+
+
++
Sodium poly(acrylate) absorbs 500 times its own mass of
water.
-
-
-
-
-
Predicting solubility from
solute and solvent polarities
Learning intention
Learn how the polarity of both
the solute and solvent molecules
influences solubility.
Solvent Action
Solvent Action
A liquid that a substance dissolves in is called a SOLVENT.
Solvents can be either polar or non-polar molecules.
Immiscible liquids do not mix, e.g. oil and water, however,
non-polar liquids are miscible with each other.
Polar solvents will usually dissolve polar molecules.
Non-polar solvents will usually dissolve non-polar molecules.
Water is a polar molecule so it is a polar solvent.
d+
H
Water has a polar covalent bonding
between O and H.
H
d+
d+
O d-
dd+
Dissolving in Water
Ionic Compound
dissolving
in water:
d-
d+
d+
-
d+
+
-
d+
+
-
d+
d- + ddd+
d+
+
d+
d-
d+
d+
Hydrated
ions
dd+
d+
d-
d+
-
d+ d+
d+
d-
Dissolving in Water
Dissolving in Water
Pure Hydrogen chloride is polar covalent. When water is
added it breaks to produce ions
dd+
H
d-
Cl
d+
d+
dd+
d+
d+
d+
d- H+ ddd+
d-
d+
d+
d+
Hydrated
ions
d+
d+
d-
d+
Cl-
d+ d+
d+
d-
Dissolving in Water
Generally, covalent molecules are insoluble in water.
However, small moleculeslike ethanol (C2H5OH),
with a polar O-H functional group, will dissolve,
dd+
H2O
d+
Ethanol
H
H
H
C
C
H
H
d-
O
H d+
Bond Strengths
Bond Type
Strength (kJ mol –1)
Metallic
80 to 600
Ionic
100 to 500
Covalent
100 to 500
Hydrogen
40
Dipole-Dipole
30
London’s forces
1 to 20