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!