cfe higher chemistry unit 1 - chemical changes and structures · 2019. 12. 20. · oxalic acid...
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December 15 1
Name ___________________________________ Class _____________
CfE Higher Chemistry
Unit 1 - Chemical Changes and
Structures
Key Area page
(1) Controlling the Rate 2 - 10
(2) Reaction Profiles 12 - 18
(3) Periodicity 19 - 29
(4) Bonding, Structure and Properties 30 – 50
December 15 2
(1) CONTROLLING THE RATE
In industry new products, such as medicines, ammonia, paint and soap powder must be produced
in large or small quantities quickly and cheaply. Chemical engineers must determine the ideal
reaction conditions before mass production begins. If reaction rates are too low (too slow)
profit margins will be too small; if too high production may be unsafe leading to the risk of a
thermal explosion. Extractor fans remove flour dust from the air in large bakeries to prevent
the risk of thermal explosion. In N5 you learned 4 factors which affect the rate of a reaction:
1. _______________ 2. ________________ 3. ______________ 4. ________________
All of these factors can be explained by Collision Theory. This theory states that before a
reaction can take place, the particles must collide with each other.
We are surrounded by air containing 78% Nitrogen and 21% Oxygen gas molecules which are
colliding all the time, yet we are not poisoned by dangerous brown nitrogen dioxide gas!!
Draw electron cloud diagrams of oxygen and nitrogen below:
Oxygen molecule Nitrogen molecule
When these molecules move towards each other, it is their electron clouds which first collide.
Think about this!!!
Electron clouds contain only negative charges so as they collide they must _________ each
other. The molecules do not move fast enough to overcome this repulsion. These collisions are
not successful. Energy, eg. lightning, needs to be supplied.
The minimum quantity of energy needed to start a reaction is called the Activation Energy
(EA). The units are kilojoules per mole (kJmol-1). Oxygen and nitrogen do not react quickly at
room temperature (RT) so the activation energy for this reaction must be __________.
The reaction between hydrogen and oxygen does not take place at RT even though it is an
explosive reaction. This reaction also has a _______ activation energy. Energy, in the form
of a burning splint is used to test for the hydrogen ‘pop’.
Neutralisation reactions take place very quickly at RT, so they must have very _______
activation energy:
H+(aq) + OH-(aq) H2O(l)
December 15 3
In photography light provides the activation energy for silver salts on photographic film to be
reduced to silver atoms:
Ag+ + e− Ag(s)
Ultra violet light provides the activation energy for the reaction between hydrogen and
chlorine to produce hydrogen chloride gas:
H2 + Cl2 2HCl
We will now examine how factors make collisions successful and increase reaction rates.
Collision Theory and Concentration
Concentration is a measure of the number of particles of reactants in a known volume. As
concentration increases the number of collisions ________________ so the rate of the
reaction __________________.
Collision Theory and Particle Size (Surface Area)
As particle size ______________ the number of collisions ________________ so the rate
of the reaction __________________.
OR
As surface area ______________ the number of collisions ________________ so the rate
of the reaction __________________.
December 15 4
Collision Theory and Temperature
The temperature of a substance is a measure of the average kinetic energy (KE) of all of its
particles. At any given temperature all the particles in a substance move at different speeds;
some are speeding up and some are slowing down.
At 200C particles all the particles in a substance move at different speeds so there is a
distribution of KE; some particles lose energy (slow down) some particles are gain energy (move
faster). So the overall pattern always stay the same!! This is shown in the graph below:
The graph shows that a very small number of particles have very low KE and a very small
number of particles have very high KE. The majority of particles have an _______________
KE because they are either losing or gaining energy.
The activation energy (EA) in the minimum quantity of energy needed by particles to collide
successfully and react. The shaded area represents the number of particles which have KE
equal to, or greater than, the activation energy (EA).
If the temperature is increased from T1 to T2 (200C to 300C) all the particles move faster
(have greater KE) so more particles have the minimum quantity of energy (EA) to collide and
react. The whole distribution curve shifts to the right:
The 2 graphs can be superimposed to show the effect of increasing the temperature:
December 15 5
The graph now has a larger shaded area which represents the increased number of particles
with energy equal to, or greater than, the activation energy.
Effect of Temperature on Reaction Rate
As temperature increases, particles move ____________ and collide with more
___________, so the rate of ___________________ collisions increases therefore
reaction rate ___________________.
A small rise in temperature can result in a large number of particles having energy equal to, or
greater than, the activation energy. This is why reaction rate approximately doubles with every
100C increase in temperature. This is shown in the graph below:
Most chemical reactions can be represented by this graph but there are other possibilities
which you are already familiar with, for example, explosions and enzyme reactions.
Demonstration Methane (or Petrol + Oxygen) (or Test for Hydrogen)
December 15 6
Although most chemical reactions follow the above pattern, there are other possibilities:
Catalysts and Activation Energy
In N5 we learned that catalysts speed up chemical reactions without getting used up in the
reaction. In fact, catalysts increase reaction rate by lowering the Activation Energy for the
reaction.
Draw an energy distribution graph below to show this:
Collision Geometry
The balanced equation for the reaction of methane
with oxygen is:
_____________________________________
The graph shows that, as the temperature
increases, the rate of an explosive reaction
increases ________________________!
An example of an enzyme reaction is the hydrolysis
of starch into glucose by the action of the enzyme,
amylase:
_____________________________________
As the temperature increases, the rate of an
enzyme reaction _______________________
then _____________________ as the enzyme is
________________________.
Using a catalyst lowers the Activation
Energy so __________ particles have
Kinetic Energy equal to, or greater than,
the Activation Energy.
This results in an __________________
in the number of __________________
collisions so reaction rate
______________.
December 15 7
Particles must also have the correct orientation with each other so that they can react.
Consider the reaction:
CO(g) + NO2(g) CO2(g) + NO(g)
correct collision geometry successful collision
repulsion unsuccessful collision
So there are 2 conditions which must be met for a successful collision to occur:
Correct collision geometry
Minimum kinetic energy
Effect of Pressure on Reaction Rate – ONLY APPLIES TO GASES!!
We will study pressure in more detail in unit 3 but in the meantime, this is very similar to the
effect of increasing the concentration of particles.
An increase in pressure of a mixture of gases increases the number of collisions and if particles
are have energy equal to, or greater than, the Activation Energy, AND the correct collision
geometry, the rate of the reaction will increase.
Methods of Monitoring Reaction Rate
December 15 8
(Revision – N5 Average rate Examples and Rate Graphs Sheet)
There are many methods of monitoring the course of a chemical reaction such as measuring:
Temperature
Volume of a gas produced
pH
Colour
Concentration
It is very difficult to monitor the course of a chemical reaction by measuring the change in
concentration of a reactant or product over a period of time. However, reactions can be
monitored by measuring another property which is related to concentration.
At the beginning of a reaction there is a __________ concentration of reactants and a
___________ concentration of products.
As the reaction proceeds the concentration of reactants ______________________ while
the concentration of products ______________________.
Relative rate of Reaction (You may carry out a practice Outcome 1 if time allows)
This just means we will compare reactions carried out at different concentrations or
temperatures etc. to find out how this affects the reaction rate. We will measure the time (t)
it takes to produce the same quantity of product in each reaction. Since the same quantity of
product is always produced we can say the quantity is 1.
quantity of product = 1
time is in seconds = t
Relative Rate = 1
t (s)
The unit of relative rate is ________.
Effect of Concentration on Relative Rate (Outcome 1)
December 15 9
It is important to try to understand the chemistry of the reaction first. We start with a
balanced chemical equation:
H2O2(aq) + 2H+(aq) + 2I-(aq) 2H2O(l) + I2(aq)
This equation tells us that: hydrogen peroxide (H2O2) reacts with potassium iodide to form
water and iodine in acidic conditions – how do we know that conditions are acidic?
Iodide ions (I-) are supplied using a potassium iodide solution. Potassium ions (K+) are not in the
equation. Why not?
The aim of the experiment is to find out the effect of changing the concentration of
potassium iodide solution on the rate of reaction with hydrogen peroxide solution.
If we are given one bottle of potassium iodide solution with a concentration of 0.1moll-1, how
do we change the concentration and keep this a fair test?
25
How do we know the reaction is finished?
H2O2(aq) + 2H+(aq) + 2I-(aq) 2H2O(l) + I2(aq)
We need a 2nd chemical reaction and an indicator to ‘see’ when the reaction is complete.
Thiosulfate ions in sodium thiosulfate react with the iodine molecules as soon as they are
produced so the solution will be remain colourless.
2S2O32-(aq) + I2(aq) S4O6
2-(aq) + 2I-(aq)
We will add the same quantity of thiosulfate ions to every experiment until they are all used
up. This quantity is 1. The time, in seconds (s), for all the thiosulfate to be used up is t. So the
rate of the reaction is 1/t (s-1). Changing the concentration of KI (aq) will affect the rate of
the reaction.
When the thiosulfate ions have been used up, iodine molecules, which are still being produced,
can be detected with starch which acts as an indicator. We will see the colour change:
_______________ _________________
See experiment workcard (PPA 1) for instructions.
December 15 10
Read all instructions and safety before you start.
Ask your teacher if you have to write up Outcome 1.
Conclusion
As the concentration of a reactant increases, the reaction rate ____________________, ie.
rate is proportional to concentration.
Insert the graph you obtained below as a reminder of the investigation.
Effect of Temperature on Relative Rate (Outcome 1)
Oxalic Acid reacts with an acidified solution of potassium permanganate:
5(COOH)2(aq) + 6H+(aq) + 2MnO4(aq) 2Mn2+(aq) + 10CO2(g) + H2O(l)
The course of the reaction can be followed by measuring the time it takes for the reaction
mixture to turn colourless.
Experiment - See experiment work card (PPA 2) for instructions.
Read all instructions and safety before you start.
Ask your teacher if you have to write up Outcome 1.
Conclusion
December 15 11
As the temperature of a reaction increases, the reaction rate ____________________, ie.
rate is proportional to concentration.
What temperature rise is needed to double the reaction rate? _____________.
A _______ increase in temperature _________ change in reaction rate.
Insert the graph you obtained below as a reminder of the investigation
(2) REACTION PROFILES
December 15 12
Exothermic Reactions
Most chemical reactions are exothermic releasing energy, usually in the form of heat, to the
surroundings (container, air etc) so there is a temperature rise.
Experiment Demo: Very Exothermic Reactions!
The Pathway of an Exothermic Chemical Reaction
When reactant particles collide successfully bonds in the reactant particles are broken. Energy
is needed to break these bonds. This is the Activation Energy (EA).
Energy is released when bonds are made in the new products. The overall energy change
depends on the quantities of energy involved in the bond breaking and bond making steps.
If the energy needed to break the bonds in the reactants is _________ than the energy
released when the bonds in the products are made then the reaction is exothermic.
We can show this is in an energy profile diagram:
In an exothermic reaction the products have less energy than the reactants so heat energy is
gained by the surroundings – there is a temperature rise. All neutralisation reactions are
exothermic.
Endothermic Reactions
December 15 13
Reactions in which energy is absorbed from the surroundings are called endothermic reaction
eg. the reaction between barium hydroxide pentahydrate and ammonium thiocynate..
1. Add 3 heaped spatula of powdered barium hydroxide to a boiling tube.
2. Measure and note the temperature of this solid.
3. Add 3 heaped spatula of ammonium thiocynate to the barium hydroxide and use the
thermometer to mix the solids.
4. Measure and record the lowest temperature reached. Touch the boiling tube as well.
Results
Temperature of barium hydroxide _______ oC
Temperature of reaction mixture _______ oC
Temperature decrease ________ oC
This is an ________________________ reaction.
The Pathway of an Endothermic Chemical Reaction
Complete the energy profile diagram below.
The energy contained in the products is _________ than the energy in the reactants. Energy
in the form of ________ must have been _________ by the surroundings.
Enthalpy Change
December 15 14
The net energy change in a chemical reaction is normally referred to as the Enthalpy Change
for the reaction and is given the symbol _______.
The enthalpy change is calculated using the equation: ΔH = HP - HR
ΔH = _________________________
HP = __________________________
HR = __________________________
It is impossible to measure the absolute enthalpy of a chemical however it is fairly easy to
carry out a chemical reaction and measure the resulting enthalpy changes.
The enthalpy change for an exothermic reaction will have a ___________________ sign
which must be shown eg. ΔH = -550kJ mol-1
The enthalpy change for an endothermic reaction will have a _______________ sign eg.
ΔH = +230kJ mol-1
Activation energy barriers
December 15 15
In all chemical reactions, the Activation Energy (EA) is the minimum kinetic energy required by
colliding particles before a reaction will occur. Sometimes this energy can be supplied in the
form of heat eg. when using a Bunsen burner.
The unit of Activation Energy is _______________.
This information can be added to the diagrams you completed earlier:
Catalysts and Activation Energy
December 15 16
Catalysts _____________ the rate of a chemical reaction by providing an alternative reaction
pathway with a _____________ activation energy.
Add this information to the diagrams below:
From both diagrams it is clear that using a catalyst ___________________________ on the
overall enthalpy change.
Work through Potential Energy Examples How Catalysts Work
December 15 17
This experiment below confirms that a catalyst can ‘take part’ in a reactions without being used
up. Cobalt ions are used to catalyse the reaction between potassium sodium tartrate solution
and hydrogen peroxide solution (a homogeneous catalyst). Cobalt ions will be regenerated.
Rochelle Salt Experiment
1. Add about 2cm3 potassium sodium tartrate solution to a boiling tube.
2. Add a few drops of cobalt (II) chloride solution until the mixture is a distinct pink colour.
3. Gently warm the mixture in a Bunsen flame – do not boil!!
4. Add a few drops of hydrogen peroxide to the heated mixture and note all colour changes.
5. Add a few more drops of hydrogen peroxide and note any further changes.
6. Draw diagrams here.
Co2+(aq) Co3+(aq) Co2+ (aq)
Catalysts can be briefly chemically changed during the reaction and regenerated at the end of
the reaction ie. it is not used up:
Mechanism
A catalyst provides an alternative pathway, with a lower activation energy, for a chemical
reaction. So less energy is needed to start the reaction. Many catalysts are solid
(heterogeneous) and simply provide a surface upon which the chemical reaction takes place at
active sites. Add labels to the diagram below:
Heterogenous Catalysts
1. Add 2cm3 hydrogen peroxide to a test tube and test for the production of oxygen.
December 15 18
2. Add a pinch of Manganese(IV) Oxide powder.
3. Note the difference in reaction rate and test for oxygen production.
Activated Complex
For successful collisions reactant particles must collide with energy equal to, or greater than
the activation energy of the reaction. In the example below the diatomic molecule XY
decomposes into its elements as follows:
2XY(g) X2 + Y2
As the molecules collide, the X – Y bonds are weakened and partial bonds are set up between
the X atoms and between the Y atoms – this is the activated complex.
Reactants ‘Activated Complex’ Products
The activated complex is a very unstable arrangement of atoms at an intermediate stage
between reactants and the products. This occurs at the top of the activation energy barrier
and has a very high potential energy:
The activated complex has a fleeting existence; there is an equal chance that the complex
could lose energy to reform the reactants or retain enough energy to become new products.
An example of this is shown below in the reaction between ethene and bromine:
ENZYMES are BIOLOGICAL CATALYSTS. These are mentioned more in Unit 2.
3) PERIODICITY
The Modern Periodic Table
December 15 19
The Russian chemist, Dimitri Mendeleev (1839 – 1907), arranged elements in order of
increasing atomic __________. He also produced columns of elements with similar
_______________ properties. Mendeleev left gaps for elements yet to be
______________. He made predictions about undiscovered elements; he correctly predicted
the properties of __________________, an element he called eka-silicon. He predicted this
would fill the gap between ________________ and ______________ in the Periodic Table.
The current Periodic Table is based on the one drawn by Mendeleev. Each element in a period
has an atomic _______________ which increases by one across the period. This is due to the
difference in the number of ________________ in the nucleus of successive elements.
Elements in the same group have the same number of _______________ in the outer energy
level (shell): alkali metals have ____ electron in the outer energy level; halogens have ____
electrons in the outer energy level. ______________ gases are a group of unreactive gases.
Their lack of reactivity is due to these elements all have _________ outer energy levels.
Structure of the first 20 elements (you must learn this)
Elements with atomic numbers 1 – 20 in the Periodic Table can be grouped according to their
bonding and structure:
All atoms or molecules are held together by weak interatomic or intermolecular forces
called van der Waals forces.
There are several categories of van der Waals forces which we will find out about in this
section of work.
Covalent Molecular Gases
The simplest of are the diatomic elements eg. hydrogen, nitrogen, oxygen and the halogens:
Metals consist of a giant lattice of positively charged
‘ions’ and delocalised outer electrons. Delocalised
electrons allow metals to _____________
__________________. The attraction of ‘ions’ and
outer electrons is called metallic bonding. Metallic
bonds are very strong: 80 – 600kJmol-1. This means
that up to 600kJ of energy is needed to break 1 mol
of metallic bonds. The greater the number of
delocalised electrons, the greater the charge on the
metal ‘ions’, the greater the strength of the metallic
bonds.
Monatomic noble gases exist as separate or
discrete atoms.
Noble gases can be cooled down to the liquid or
solid state.
What holds these atoms together when cooled
down?
December 15 20
Each molecule contains 2 atoms held together by very strong covalent bonds. Covalent bonds
are approximately as strong as either metallic or ionic bonds.
Hydrogen, nitrogen, oxygen, fluorine and chlorine are all gases. These gases can be cooled to
form liquids and solids.
Melting and boiling point information in the data book confirm that bromine is a liquid and iodine
a solid.
What holds all of these diatomic molecules together in the liquid and solid states naturally or
when cooled? _______________________
Covalent Molecular Solids
Phosphorus is made up of discrete
molecules although it is a solid. Each
molecule consists of 4 phosphorus atoms
held together by covalent bonds. All of
the P4 molecules are held together by
weak van der Waals forces of
attraction.
Sulfur is made up of discrete molecules
even though it is a solid. Each molecule
consists of 8 sulfur atoms held together
by covalent bonds. All of the S8 molecules
are held together by weak van der Waals
forces of attraction.
Buckminster fullerenes, known as fullerenes,
were discovered in 1985.
Fullerenes are discrete molecules containing 60
or more atoms held together by covalent bonds.
Molecules are held together by weak van der
December 15 21
C60 molecules have a spherical shape
(almost identical to a football). Other shapes such as nanotubes
can also be made.
Fullerenes have higher melting points than sulfur or phosphorus. Explain this:
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
Which has a higher melting point sulfur or phosphorus? __________________
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
Covalent Network Solids
December 15 22
Carbon exists naturally in two forms, diamond and graphite, both are covalent networks:
B12 molecular units interlock to
form a 3 dimensional covalent
network structure. All bonds are
covalent.
Although the word molecule
appears in this description you
must learn that this is not the
same as the discrete molecular
solids on the previous page.
In diamond, each carbon atom is at the centre of
a regular tetrahedron and is surrounded by four
other carbon atoms at the corners of the
tetrahedron.
All bonds are covalent; therefore diamond has
very ________ mp and bp.
Diamond is the hardest substance known to man
(and woman).
In graphite, each carbon atom forms covalent
bonds with only 3 neighbouring carbon atoms to
form layers (like graphene).
The layers are held together by weak van der
Waals forces of attraction.
Every time we use a pencil, layers of graphite
slide on to the paper.
Graphite is also use as a lubricant between
moving metal parts.
Silicon has a 3-dimensional
covalent network lattice
structure similar to diamond so
also has a very ______ mp.
December 15 23
There are 3 polymorphs (different forms) of carbon. They are:
______________________ and __________________________(covalent networks)
______________________ (discrete molecules)
Complete the following summary table using the appropriate type of bonding and structure.
A
B
C(i)
C(ii)
D
Explain why covalent network elements have high melting and boiling points.
December 15 24
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
Explain why discrete monatomic elements and discrete molecules have low melting and boiling
points.
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
Explain why graphite conducts electricity and why diamond does not.
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
Explain why diamond is hard and why graphite is not.
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
How does the structure of diamond and graphite differ from fullerenes?
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
Trends in the Periodic Table
December 15 25
Data book information shows that there are trends in the periodic table across a period and
down a group. We will examine:
Covalent Radius (used as a measure of atomic size)
Electronegativity
Ionisation energy
Covalent Radius
The covalent radius is about half the distance between the two nuclei of covalently bonded
atoms of an element.
Trend
Across a period covalent radius ___________________.
Down a group covalent radius _____________________.
Explanations
Across a period the number of protons ________________ which increases
_______________ charge. Increased nuclear charge attracts electrons closer to the nucleus
so covalent radius ___________________.
December 15 26
Down a group the number of filled electron shells increases. So even though there are more
protons and an increased nuclear charge, the shielding effect of the inner electron shells
prevents outer electrons being strongly attracted to the nucleus.
Draw a graph of atomic size vs atomic number for elements 3 to 20
Draw dotted lines between atomic number 9 and 11 then between 17 and 19
Draw lines to ‘join the remaining dots’
Add a page number and file your graph here!
Use your graph to answer the following questions stating the trend and the explanation.
How does the atomic size of lithium compare with that of fluorine? Explain.
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
How does the atomic size of lithium compare with that of caesium? Explain
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
Atomic size is a periodic property, ie. a pattern is repeated across each period with
elements in the same groups occurring at the same positions on the ‘waves’.
Electronegativity
December 15 27
Electronegativity is a measure of an atom’s nuclear attraction for the electrons in a covalent
bond. Electronegativity values (Pauling Scale) are found in the data book.
Trend
Across a period electronegativity values ___________________.
Down a group electronegativity values ___________________.
Explanations
Across a period the number of protons ________________ which increases
_______________ charge. Increased nuclear charge attracts/pulls bonding electrons closer
to the nucleus.
Down a group the number of filled electron shells increases. So even though there are more
protons and an increased nuclear charge, the shielding effect of the inner electron shells
prevents bonding electrons being strongly attracted to the nucleus.
Draw a graph of electronegativity values against atomic number for the first twenty
elements
Use a dotted line between the noble gases and the elements in Group 1
‘Join dots’ for elements in same period
Add a page number and file your graph here
Use your graph to complete the following statements
Electronegativity can be described as a periodic property because ___________________
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
There are no electronegativity values for the noble gases because ____________________
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
Ionisation energy
December 15 28
The first ionisation energy of an element is the energy required to remove one electron from
each atom in one mole of gaseous atoms of the element. Energy is needed for this process so
It is an _______thermic process:
K (g) K+ (g) + e- ΔH = + 425kJ mol-1
Trend
Across a period first ionisation energy values ___________________.
Down a group first ionisation energy values ___________________.
Explanations
Across a period the number of protons ________________ which increases
_______________ charge. Increased nuclear charge attracts outer electrons closer to the
nucleus so more energy is needed to remove an outer electron.
Down a group the number of filled electron shells increases. So even though there are more
protons and an increased nuclear charge, the shielding effect of the inner electron shells
prevents strong attraction of outer electrons so less energy is needed to remove an outer
electron.
2nd, 3rd and 4th Ionisation energies
More than one electron can be removed from an atom. The 2nd ionisation energy of an element
is the energy required to remove one electron from each ion in one mole of gaseous ions of
the element. It is also an endothermic process.
The calculation below shows the total energy input required to convert one mole of Magnesium
atoms into one mole of Magnesium Mg2+ ions:
Calculate the energy required to convert one mole of Al atoms into one mole of Al3+ ions:
Why is the 2nd ionisation energy of sodium so much bigger than the first ionisation energy?
December 15 29
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
Draw a graph of first ionisation energy against atomic number for the first twenty
elements
Use a dotted line between the noble gases and the elements in Group 1
‘Join dots’ for elements in same period
Add a page number and file your graph here
Use your graph to complete the following statements
The first ionisation energy can be described as a periodic property because _________
_____________________________________________________________________
_____________________________________________________________________
The first ionisation energy of chlorine is _________________ than that of sodium because
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
The first ionisation energy of fluorine is _________________ than that of iodine because
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
The trends, and reasons for the trends, for covalent radius, electronegativity values and
ionisation energies are very similar.
c) Periodic trends in ionisation energies and covalent radii - Trends in the Periodic Table and
Bonding - Welcome to the NQ Higher Sciences website
d) Periodic trends in electronegativity - Trends in the Periodic Table and Bonding - Welcome to
the NQ Higher Sciences website
(4) Bonding, Structure and Properties
December 15 30
Ionic Compounds
Compounds formed between metals and non-metals are usually, but not always ionic. Ionic
compounds are formed if metal and non-metal atoms have large differences in electronegativity
values. All ionic compounds are solid at room temperature so must have high melting points. If
not, the compounds are covalent. Melting and boiling point information needs to be examined
to determine the type of bonding.
Metals react with non-metals to form ionic compounds. A transfer of electrons from one atom
to another results in metal atoms ____________ electrons to form an ion with a
____________________ charge. Non – metal atoms ____________ electrons to form an
ion with a ____________________ charge. These ions have the stable electron arrangement
of a noble gas. Oppositely charged ions ________________ each other. This electrostatic
attraction is known as an ______________ bond.
f) Bonding continuum - Trends in the Periodic Table and Bonding - Welcome to the NQ Higher Sciences website
All ionic compounds have ________ melting points therefore are all _________ at room
temperature. All conduct electricity when _____________ or in _________________
because ions are _________________. They do not conduct in the ______________ state.
Ionic bonds are very strong: up to 600 kJmol-1.This means that up to 600kJ of energy is
needed to break 1 mol of ionic bonds. The strength of ionic bonds depends on the charge on
each ion; the higher the charge, the stronger the bond.
Covalent Bonds - revision
Non – metal atoms join together by covalent bonds. A covalent bond is a ______________
_________ ___ ______________. Each atom has the stable electron arrangement of a
______________ gas.
Draw diagrams of methane and ammonia molecules on the next page to show:
Overlapping electron clouds
Shape of the molecules
In a crystal lattice eg. NaCl (or Na+Cl-), the
formula tells us the ratio of ions present in
the lattice.
The formula does not tell us the actual
number of ions present.
December 15 31
Methane overlapping clouds Ammonia overlapping clouds
Methane shape Ammonia shape
Discrete Molecules
Discrete molecules have very strong internal covalent bonds but only weak intermolecular
bonds. Use mp and bp information from the data book to complete the table.
Name of Formula Structure State at Room Temperature Carbon
dioxide
Ammonia
Hydrogen
oxide
Hydrogen
chloride
Hexane
Discrete covalent molecular compounds are either _____________ or ___________.
December 15 32
Solid Covalent Network Compounds
Silicon dioxide (_________) commonly known as sand contains only _________________
bonds in its network so has a very high ________ and a structure similar to diamond:
Silicon carbide (________) commonly known as carborundum contains only _______________
bonds in its network so has a very high ________ .
Both silicon dioxide and silicon carbide are very hard. Silicon carbide is used as an abrasive for
cutting and grinding the surfaces of tools.
Summary table for covalent structures
Type of Bonding and Structure Properties
Covalent network solids
_________________ melting points
_______________ of electricity (carbon in the
form of ________________ is the exception)
Covalent molecular solids
_________________ melting points
_______________ of electricity
Covalent molecular gases and liquids _________________ melting points
_______________ of electricity
December 15 33
Types of Covalent Bonds
Covalent bonds are very strong: up to 600kJmol-1. This means that up to 600kJ of energy is
needed to break 1 mol of covalent bonds. Electrons in the covalent bonds are held in place by
the force of attraction of positive nuclei. Show these attractions in the diagram below:
Electronegativity values tell us which nucleus has a greater attraction for bonding electrons.
If there are 2 atoms of the same element, electronegativity values are equal. In a molecule of
hydrogen, each atom has an electronegativity value of ______ so each nucleus has an equal
attraction for the bonding electrons. This means electrons are more likely to be found in the
centre of the bond. This is a pure covalent bond. All diatomic elements have pure covalent
bonds.
Electrons are constantly moving (wobbling) so they could be found nearer the nucleus of one of
the atoms than the other. This means that one end of the molecule could become negatively
charged while the other end becomes positive. This is known as a dipole.
Electrons return to the centre to become pure covalent again before moving to the opposite
end of the molecule so the polarity changes again. All diatomic elements contain pure covalent
bonds which exhibit temporary dipoles. This is shown using the Greek symbol delta which
means very small ________.
Show the changes in polarity for the elements chlorine and iodine the spaces below:
December 15 34
Comparing the electronegativity values of atoms in diatomic molecules made up of different
elements, produces a different outcome. For example Hydrogen Iodide:
electronegativity values: hydrogen = ______ iodine = _______
The nucleus of ______________ has a greater attraction for bonding electrons than
________________. Electrons will always be nearer ______________. This results in a
polar covalent bond because each end (pole) has a different charge. The molecule has a
permanent dipole:
Now repeat the process for Hydrogen Chloride and then Hydrogen Fluoride:
Which of the 3 molecules above has the most polar covalent bond? ___________________
Why? ________________________________________________________________
Summary so Far
Pure covalent bonds occur in diatomic elements because the atoms have identical
electronegativity values. Each nucleus has the same attraction for bonding electrons.
Temporary dipoles occur. Polar covalent bonds occur in diatomic molecules because the atoms
have different electronegativity values. One nucleus has a greater attraction for bonding
electrons. Permanent dipoles occur due to unequal distribution of internal charges.
e) Polar covalent bonds - Trends in the Periodic Table and Bonding - Welcome to the NQ Higher Sciences
website
The Bonding Continuum
Pure covalent bonding and ionic bonding are two extremes of polar bonding. Atoms with
identical electronegativity values (pure covalent bonds) are found at one end of the spectrum.
Further along polar bonds with increasing differences in electronegativity values are found.
Ions are formed if the difference in electronegativity is such that the atom with the greatest
attraction for bonding electrons completely pulls electron(s) away from the other atom. Ionic
bonding appears at the furthest point in the continuum.
f) Bonding continuum - Trends in the Periodic Table and Bonding - Welcome to the NQ Higher Sciences website
The bonding continuum:
_____________________________________________________________________
December 15 35
Polar and Non-Polar Molecules
Molecules can contain bonds with permanent dipoles (polar covalent bonds) but overall the
molecule can be non-polar.
It is the symmetry of the charge distribution which determines the overall polarity of the
molecule.
Methane and Butane
By substituting just one atom in each of these molecules, we immediately affect the
distribution of charge around the carbon atom.
C
H
H
H H
C
H
F
H H
Each bond in the methane molecule is
___________ because the electronegativity
values are _____________________.
The bonds are all arranged so that the charge is
symmetrical around the carbon atom. The
outside of the molecule is uniformly - in all
directions so overall the molecule is non-polar.
Hydrocarbon molecules, like butane, are usually
non-polar because the spatial arrangement of
charge is distributed symmetrically.
Build a molecule of butane to see this for
yourself.
The carbon – halogen bond is very polar because of
the difference in electronegativity values.
Each molecule has lost its symmetry of charge so the
molecule is polar.
Show the charges using the delta symbol to see
how each molecule is affected and build the
molecule too!
December 15 36
Look at the molecules below and put the correct charge beside each delta symbol.
Decide whether each molecule is polar or non-polar by looking at the symmetry of
charge.
Build the molecules to help you ‘see’ this.
_________________ ________________ _________________
__________________
__________________ _________________
December 15 37
_____________________
Activity
The polar nature of some molecules can be ‘seen’ in an electric field by holding a ‘charged’
plastic rod close to a stream of the test liquid.
Draw the structures of the molecules in the table and predict their polarity before carrying
out the activity:
Name of Liquid Polarity prediction Actual Polarity
Water
Hexane
ethanol
propanone
Experiment
December 15 38
Deflection will occur only if the molecules in the liquid are ______________.
Make sure you could reproduce the diagram in an exam and explain it.
Intermolecular Forces
Theoretically every element can exist in 3 states: solid, liquid and gas. When cooled sufficiently
elements such as the noble gases, diatomic chlorine gas and liquid bromine can become liquid
and/or solid.
What holds these atoms or molecules together in the liquid and solid states?
Intermolecular forces between atoms or molecules are known as van der Waals forces of
attraction. There are different types of van der Waals forces depending on the
electronegativity values of atoms being attracted but they all occur in exactly the same way.
London dispersion forces
Noble gases are m________________. When cooled atoms lose energy and move closer
together. All electrons are constantly moving so atoms can exhibit temporary dipoles.
Add partial charge delta symbols to the noble gas atoms below to show this:
December 15 39
Only the outer electrons are shown to illustrate temporary dipoles but all electrons can move
to contribute to the strength of the partial charge. The more electrons in the atom the
stronger the partial charge.
Opposite charges attract to form van der Waals forces called London dispersion forces,
which hold the atoms together in the liquid or solid state:
Since electrons are constantly moving the London dispersion forces will continually break and
reform as poles change charge.
animations - Welcome to the NQ Higher Sciences website
London dispersion forces are extremely weak and easy to break but their strength varies
depending on the number of electrons in the atom. The more electrons in the atom the
stronger the partial charge – so the London dispersion forces will be stronger.
Melting and boiling point information confirms this. In general, the melting and boiling points
of noble gases are ________________________ so the London dispersion forces holding the
atoms together are very __________ and __________ to break.
December 15 40
London dispersions forces
The simplest are the diatomic elements, hydrogen, nitrogen, oxygen and the halogens. Each
molecule contains only 2 atoms with the same electronegativity value resulting in
__________________ dipoles within the molecules (intramolecular). Weak London dispersion
forces hold molecules together in the liquid and solid states when cooled sufficiently.
Draw dotted lines to show all possible London dispersion forces between molecules forces
Hydrogen, nitrogen, oxygen, fluorine and chlorine are gases at RT but bromine is a liquid and
iodine is a solid.
Why does the state change as we move down the halogens? (Open ended question?)
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
____________________________________________________________________
Which molecular element should have the highest melting point sulphur or phosphorus and
why?__________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
London dispersion forces are extremely weak and easy to break but their strength depends
on the number of electrons in the molecule. The more electrons in the molecule the stronger
the partial charges – so the stronger the London dispersion forces and harder to break.
Permanent dipole/permanent dipole interactions
Additional van der Waals forces occur along with the London dispersion forces when
molecules are polar. Permanent dipoles occur when atoms have different electronegativity
values – the bigger the difference the stronger the partial charges.
December 15 41
The additional intermolecular forces which occur are known as permanent dipole/permanent
dipole interactions. These are stronger than London dispersion forces. animations - Welcome
to the NQ Higher Sciences website
The permanent dipoles in hydrogen chloride can be shown as follows:
Draw dotted lines to show all possible permanent dipole/permanent dipole intermolecular
forces
The permanent dipoles in hydrogen iodide can be shown as follows:
Draw dotted lines to show all possible permanent dipole/permanent dipole intermolecular
forces in HI
Which compound has the higher melting point, HCl or HI ? _________________________
December 15 42
Why? _________________________________________________________________
_____________________________________________________________________
Hydrogen bonding
Hydrogen bonding is the strongest of all the van der Waals forces. They also co-exist with
London dispersion forces. Hydrogen bonding is stronger example of permanent
dipole/permanent dipole interactions. Hydrogen bonds only occur when a hydrogen atom in the
molecule is covalently joined to one of the three most electronegative elements: nitrogen,
oxygen or fluorine. animations - Welcome to the NQ Higher Sciences website
The permanent dipoles in hydrogen fluoride can be shown as follows:
Draw dotted lines to show all possible hydrogen bonds HF.
More Examples
The permanent dipoles in nitrogen hydride can be shown as follows:
Draw dotted lines to show all possible hydrogen bonds NH3.
December 15 43
The permanent dipoles in hydrogen oxide can be shown as follows:
Draw dotted lines to show all possible hydrogen bonds H2O.
Melting and boiling point information confirms the difference in strength between permanent
dipole-permanent dipole intermolecular forces and hydrogen bonds.
Hydrogen chloride and hydrogen iodide have __________________ mps and bps than
hydrogen fluoride, nitrogen hydride and hydrogen oxide so hydrogen bonds are
________________ than _________________________________________________.
Chemical Analysis Problem Solving Question
In the liquid state Hydrogen Fluoride can form long hydrogen bonded chains:
The large electronegativity difference between these atoms make hydrogen bonds much
stronger than other intermolecular forces (hydrogen bond energy is up to about 50 kJmol-1)
Hydrogen bonding can lead to some unusual results when analysing substances:
Mass spectrometer information on hydrogen fluoride produces 3 different masses, 20, 40 and
60 for this substance. Explain this!! (Open ended question?)
December 15 44
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
Summary of Bonding
The Importance of Electronegativity Values and the Bonding Continuum
The larger the difference in electronegativity values the more ionic character but melting and
boiling point data must also be examined.
The smaller the difference in electronegativity values the more likely a substance is polar
covalent (less ionic character).
If electronegativity values are the same the substance is considered to be pure covalent
(least/no ionic character).
Boiling Point of Polar and Non-polar compounds.
Now complete the table below for compounds and answer the questions which follow:
Substance Molecular Mass No of electrons bp oC Polar/non-polar
Butane
( )
Methanal
(HCOH)
Ethane
( )
Propanone
(CH3COCH3)
Notice that compounds with similar masses have similar number of electrons so the strength
of L_______________________________ should be similar.
How does the melting and boiling points of polar substances compare with the melting and
boiling points of non-polar substances?
Substance Type of Bonding Bond Strength/kJmol-1
Noble gases London dispersion forces (temporary) Very weak/up to 20
Molecular elements London dispersion forces (temporary Very weak/up to 20
Molecular compounds Permanent dipole-permanent dipole Weak/up to 30
Molecular compound containing
hydrogen + NOF element Hydrogen bonds Quite weak/up to 50
Covalent network Covalent Very strong/up to 600
Crystal lattice Ionic Very strong/ up to 600
Metal Metallic Very strong/ up to 600
December 15 45
_____________________________________________________________________
_____________________________________________________________________
Why is it important to compare only the melting and boiling points of substances with similar
numbers of electrons?
_____________________________________________________________________
_____________________________________________________________________
Boiling Points in Hydrides
Study the graph below. Label each line with the appropriate group number then answer
the questions which follow:
What is the General trend in boiling points for the groups in the graph?
_____________________________________________________________________
Explain this trend.
_____________________________________________________________________
_____________________________________________________________________
Now add the boiling points for the Row 2 hydrides using the information below and your data
book. Convert boiling points in OC to the Kelvin scale by adding 273.
NH3: bp = -33 OC HF: bp = 19.5 OC
Which hydride has no effect on the general trend? ______________________________
Explain why the Row 2 hydrides alter the general trend in boiling points but the hydride you
identified above does not.
December 15 46
_____________________________________________________________________
_____________________________________________________________________
The Most Important Hydride in the World ever!!!!! – Water is Weird or Wonderful
Water is a liquid (at STP). This is only possible because of hydrogen bonding between water
molecules. Hydrogen bonds are responsible for the unusually high boiling point of water.
If there were no hydrogen bonds between molecules water would exist only as a gas!!!! Think
about this!!!! Oceans and rivers would never exist and it would never rain!!!!
Humans, and other living organisms, would not exist if water was a gas because chemical
reactions in our body could not take place – an aqueous environment is needed for reactions!
When gases and liquids cool down, their atoms/molecules lose energy and move closer together
ie. density increases. This is also true for water molecules then something weird happens when
the temperature decreases to 4oC.
Between 4oC and 0oC the molecules start moving apart again – water expands when it
becomes ice!! This is because hydrogen bonds move to achieve a perfect regular crystal
structure (ice crystals). www.youtube.com/watch?v=lkl5cbfqFRM
Because ice is less dense than water (lighter) ponds and rivers freeze from the surface
downwards. This creates a layer of insulation which prevents ice forming underneath. If ice
Density increases as the
temperature falls to 4oC.
Molecules move closer together.
Between 4oC and 0oC water
molecules move apart to achieve a
perfect regular open crystal
structure.
December 15 47
was denser than water (heavier), ice would form from the bottom upwards so ponds, rivers etc.
would freeze completely, killing all fish and aquatic plants!!
Viscosity
Intermolecular bonds between molecules in liquids are continually being broken and reformed
so molecules move past each other easily. Hydrogen bonds are the strongest of the
intermolecular forces so water molecules do not move past each other as easily. Water is more
viscous.
The viscosity of different liquids can be compared by timing how long it takes a ball bearing to
fall through the liquid or a bubble to move up. The more hydrogen bonds between molecules
the more viscous the liquid.
Diagrams of Bubble Experiment and Ball Bearing Experiment (Exam question)
Complete the table below for hexane, ethanol, ethane-1,2-diol and propane-1,2,3-triol
(glycerol) in the following table to determine the number of hydrogen bonds which can be
formed by each molecule then try the experiment.
Name of
molecule Structure
Number of
H Bonds
Time for bubble to rise
(s) or ball bearing to
fall (s)
December 15 48
Conclusion
The greater the number of H bonds, the _____________ the ball bearing (or bubble) moves
through the liquid, the _________ viscous the liquid
Miscibility/ Solubility
Small polar molecules like ethanol are soluble (miscible) in water because their polar O–H
functional groups can form hydrogen bonds with the water molecules. Larger molecules, like
butanol, are relatively insoluble (immiscible) even though they have the same polar group,
because the larger non-polar hydrocarbon chain is the dominant influence.
Ethanoic acid has a formula mass of 60. Mass spectrometer information sometimes produces a
formula mass of 120 because in solution hydrogen bonds form with water molecules and also
hold the molecules together in pairs (intramolecular forces):
December 15 49
Small molecules like hydrogen chloride, HCl (g), and hydrogen iodide, HI (g), are highly
___________ because of their different _____________________ values, ie.
H---Cl H---I
These molecules are so soluble in water that when new ‘bonds’ are formed between the negative
poles of the H-I molecules and the positive poles of the water molecules enough energy is
released to break the bonds between the hydrogen and iodine atoms.
Hydrogen iodide molecules completely ionise in water to form an ______________ solution:
H+(aq) + I-(aq)
Similarly other _______________ covalent substances eg the other hydrogen halides HF(g),
HCl(g), HBr(g), and pure concentrated sulphuric acid H2SO4(l), oleic acid, all completely ionise
in water to form acidic solutions.
Non-polar Molecules
Polar compounds do not usually dissolve in non-polar solvents, like benzene or
tetrachloromethane. Benzene molecules are held together by weak London dispersion forces
between molecules in the liquid. This is also true for tetrachloromethane.
If these interactions were broken they would not release enough energy to break the stronger
permanent dipole-permanent dipole interactions between polar molecules like water.
Non polar iodine is soluble in tetrachloromethane because sufficient energy is released when
iodine molecules make new London dispersion forces with tetrachloromethane molecules to
break the existing London dispersion forces between tetrachloromethane molecules.
Solubility of Ionic Compounds
Some ionic compounds dissolve in water. The slightly negative ends of water molecules are
attracted to the __________________ ions in the crystal lattice while the positive ends are
attracted to the __________________ ions in the crystal lattice.
December 15 50
Formation of new electrostatic attractions between ions and water molecules releases enough
energy to overcome ionic bonds in the crystal lattice as shown in the diagram below:
Polar water molecules
Ionic crystal lattice Hydrated Ions
See from slide 10 onwards
f) Predicting solubility from solute and solvent polarities - Intermolecular forces - Welcome to the
NQ Higher Sciences website
Non-polar substances are the best solvents for other non-polar substances.
Polar substances are the best solvents for other polar substances and for ionic compounds.
Substances containing O-H groups are best dissolved in
water or solvents which also contain O-H groups.
In summary like dissolves like!
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