Unit 1 Chemical Changes & Structure - Duncanrig Secondary School · 2017. 6. 27. · Duncanrig Secondary School cfe Higher Chemistry Chemical Changes and Structures Page 3 of 53 Learning

Duncanrig Secondary School cfe Higher Chemistry Chemical Changes and Structures

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Duncanrig Secondary School

CfE Higher Chemistry

Unit 1

Chemical Changes &

Structure

Part 1 Controlling the Rate

Part 2 Trends in the Periodic Table

Part 3 Structure and Bonding


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Circle a face to show how much understanding you have of each statement: if you fully

understand enough to do what the outcome says, if you have some understanding of the

statement, and if you do not yet understand enough to do what the statement says. Once

you have completed this, you will be able to tell which parts of the topic that you need to

revise, by either looking at your notes again or by asking for an explanation from your teacher

or classmates.

Learning Outcomes – Controlling the Rate

By the end of this topic I will be able to:

1 Explain the effect of temperature, concentration and particle

size in terms of the energy and number of collisions (Collision

Theory).

2 State which reactions are slowest or fastest at different

points using the slope of rates graphs.

3 State that activation energy is the minimum energy required for

particles to react.

4 Draw a graph showing the effect of temperature on the kinetic

energy of particles .

5 Use activation energy on this graph to explain why higher

temperatures speed up reactions.

6 State that catalysts speed up reactions by providing an

alternative reaction pathway with lower activation energy.

7 Describe the difference between a homogeneous catalyst and a

heterogeneous catalyst.

8 Explain the adsorption, reaction and desorption stages in the

action of a heterogeneous catalyst.

9 State that catalyst poisons occupy the active site in a catalyst

and prevent it working.

10 State that enzymes are biological catalysts and give examples

of some enzymes.

11 Explain why enzymes operate at optimum temperatures and pH

values.

12 Draw potential energy diagrams for exothermic and

endothermic reactions.

13 State that enthalpy change represents the difference: ∆H =

H(products) – H(reactants).

14 State that the activated complex is an unstable arrangement of

atoms formed at the maximum of the potential energy barrier,

during a reaction.

15 Use potential energy diagrams to illustrate the effect catalysts

have on the activation energy and reaction pathway.


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Learning Outcomes – Trends in the Periodic Table


1 Define the density of an element as its mass per unit volume,

usually in gcm-3.

2 Define the covalent radius as a measure of the size of an atom

(specifically that it is half the distance between the nuclei of

two bonded atoms of an element).

3 State that the atomic size decreases across a period and

increases down a group.

4 Explain why there are changes in atomic size across a period and

down a group.

5 Define the first ionisation energy as the energy required to

remove one mole of electrons from one mole of gaseous atoms

6 Understand that the second and subsequent ionisation energies

refer to the energies required to remove further moles of

electrons.

7 Explain the trends in first ionisation energy across periods and

down groups in terms of atomic size, nuclear charge and the

screening effect due to inner shell electrons

8 Understand that atoms of different elements have different

attractions for bonding electrons.

9 Define electronegativity as a measure of the attraction an atom

involved in a bond has for the electrons of the bond.

10 State that electronegativity values increase across a period and

decrease down a group.

11 Explain the trends in electronegativity across periods and down

groups in terms of nuclear charge, covalent radius and the

presence of “screening” inner shell electrons.


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Learning Outcomes – Bonding and Structure


1

The bonding types of the first twenty elements; metallic (Li, Be,

Na, Mg, Al, K and Ca); covalent molecular (H2, N2, O2, F2, Cl2, P4,

S8 and C60 [fullerenes]); covalent network (B, C (diamond,

graphite), Si) and monatomic (noble gases)

2 Describe the bonding continuum moving from pure non-polar

covalent to ionic.

3 Explain how polar covalent bonds arise

4 Explain how van der Waals forces arise between molecules.

5 Describe what causes dispersion forces to exist between

gaseous atoms and molecules.

6 Explain how the polarity of molecules affects the strength of

dispersion forces.

7 Explain why certain molecules have a stronger type of van der

Waal force called a hydrogen bond

8 Explain how the properties of substances are affected by the

type of bonding that they exhibit..

9 Predict the solubility of a substance from information about

solute and solvent polarities.


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PART 1 CONTROLLING THE RATE

The Rate of Chemical Reactions

Everyday reactions have different speeds; some are over in a fraction of a second

(fast: like a gas explosion) while others can take years (slow: like the rusting of iron).

Most reactions occur at rates between these two extremes (medium: like a cake

baking).

Collision Theory

For a chemical reaction to occur some important things have to happen:

1. The reacting particles must collide together.

2. Collisions must have sufficient energy to produce a product.

3. The reacting particles must have the correct orientation.

Therefore anything that increases the number of and energy of collisions between

reactant particles will speed up a reaction.

Factors Affecting the Rate of a Reaction

There are three main factors affecting the rate of a chemical reaction:

a) Particle Size:

The smaller the particles, the faster

the reaction. This is because smaller

particles provide more surface area

for collision.

Example – Marble powder reacts faster with acid than marble chips.

b) Concentration:

The higher the concentration, the faster the

reaction. The higher the concentration

of solutions, the more particles you have

crowded into a small volume of liquid.

Hence, the more likely they are to

collide with each other.

Example – 2 mol l-1 hydrochloric acid reacts faster with magnesium ribbon than 1 mol l-1

hydrochloric acid.


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c) Temperature:

Although a higher temperature will cause molecules to move faster, and there may be

more collisions, this is not the main reason why higher temperature increases reaction

rate. The main reason is that more of the collisions which occur will lead to a

successful reaction. This is because at higher temperature, more particles have the

activation energy required for a reaction to happen.

As a rough guide, the rate of reaction doubles when the temperature increase by

10OC.

Example - Benedicts solution reacts faster with glucose solution at 50OC than at

25OC.

Catalysts

Even when particle size is decreased and concentration and temperature are

increased, many chemical reactions are still too slow. How can the rate of these

reactions be increased? This is especially important in today’s competitive market:

companies are constantly trying to produce more cost effective products by increasing

the rate of industrial reactions.

A catalyst is a substance which can be used to increase the rate of a chemical

reaction. The 'amount' of catalyst at the end of the reaction is the same as at the

start, i.e. the catalyst is not used up in the reaction and the catalyst can be

recovered chemically unchanged at the end of reaction. Different reactions require

different catalysts and not all reactions have a suitable catalyst.


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Collision Theory and the Activated Complex

In order to react particles must collide.

A chemical reaction will only occur if the reacting particles collide with enough kinetic

energy. The energy is required to overcome the repulsive forces between the atoms

and molecules and to start the breaking of bonds.

The minimum kinetic energy required for a reaction to occur is called the activation

energy (EA).

When the reactant particles collide with the required activation energy they form an

activated complex. This unstable intermediate breaks down.

E,g, The reaction of hydrogen and bromine

Sometimes the collisions do not result in a reaction, despite having the

minimum kinetic energy.

This is thought to be because the particles have not collided with the

correct geometry (angle) to allow the activated complex to be formed.

In the above reaction of hydrogen and bromine the particles collide side

on but if they collided end on…

H-H + Br-Br H----H-----Br----Br

no reaction occurs as the activated complex cannot be formed if only 2 of

the atoms come into contact with one another.


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Collision Theory and Concentration

The straight line graph means rate is directly proportional to the

concentrations of the reactants, i.e. double the concentration and you

double the rate. This is true of many reactions.

The faster rate is due to the increased number of collisions which must

occur with higher concentrations of reactants.

Collision Theory and Particle Size

The smaller the particle size, the faster the reaction as the total

surface area is larger so more collisions will occur.

Note

The steeper the curve the faster the reaction

The same volume of gas will be produced if the same number of

moles of reactants are used.

Concentration

(mol l-1)

Rate = 1/t

(s-1)


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The Effect of Concentration Changes on Reaction Rate


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KINETIC Theory and Temperature

Temperature is a measure of the average kinetic energy of the particles

of a substance.

At any given temperature, the particles of a substance will have a range

of kinetic energies and this can be shown on an energy distribution

graph.

NB The maximum height of T2 is always lower than T1

The graph above shows the kinetic energy distribution of the particles of

a reactant at two different temperatures.

It shows that at the higher temperature (T2), many more molecules have

energies equal to or greater than the activation energy (Ea). This

leads to an increase in the rate of successful collisions and hence reaction

rate.

A small rise in temperature can cause a large increase in the number of

particles having the activation energy and so can result in a large increase

in reaction rates.

For some reactions, a 10oC rise in temperature can double the reaction

rate.


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The Effect of Temperature Changes on Reaction Rate


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Although most chemical reactions follow this pattern there are other

possibilities.

Photochemical Reactions

Photochemical reactions are speeded up by the presence of light.

In these reactions, the light energy helps to supply the activation energy,

i.e. it increases the number of particles with energy equal to or greater

than the activation energy.

Examples of photochemical reactions are:

Photosynthesis

Alkane with bromine water

Chlorine and hydrogen gases

H2(g) + Cl2(g) 2HCl(g)

Catalysts and Reaction Rate

A catalyst is a substance which changes the speed of a chemical reaction

without being permanently changed itself.

Catalysts speed up chemical reactions by providing an alternative

reaction pathway which has a lower activation energy.

Explosive reaction Enzyme controlled reaction Enzyme reaction


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There are 2 main types of catalyst:

Homogeneous Catalysts Heterogeneous Catalysts

Homogeneous catalysts are in the same state as the reactants.

Heterogeneous catalysts are in a different state to the reactants.

e.g. Decomposition of hydrogen peroxide (solution) using manganese (IV)

oxide (solid) as a catalyst.

Manganese (IV) oxide (s)

2H2O2 (aq) 2H2O (l) + O2(g)

Common Catalysed Reactions

Catalyst Reaction catalysed Type of Catalyst

iron Haber process-ammonia manufacture Hetrogeneous

platinum Ostwald process - oxidation of

ammonia to make nitric acid

Hetrogeneous

nickel Hydrogenation of vegetable oils to

make margarine

Hetrogeneous

aluminium oxide Cracking of long chain hydrocarbons Hetrogeneous

titanium( IV) oxide Addition polymerisation -

manufacture of poly(ethene)

Hetrogeneous

Conc. sulphuric acid Esterification - making esters from

alcohol/carboxylic acid

Homogeneous

Heat

Very little

reaction

occurs

Fast reaction,

solution turns

green, gases

evolve rapidly

Solutions of

potassium

sodium tartrate

and hydrogen

peroxide (colourless)

CoCl2(aq)

+

Reaction

complete,

solution turns

pink again


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How Heterogeneous Catalysts Work

This type of catalyst is called a surface catalyst.

It works by adsorbing the reacting

molecules on to active sites and holding

them with weak bonds on its surface.

This not only causes the bonds within

the molecule to weaken but also helps

the collision geometry.

The reaction occurs on the surface with

less energy needed to form the

activated complex (lower activation

energy).

The products are formed and leave the

catalyst surface free for further

reactions

Catalyst Poisoning

A surface catalyst can be poisoned when another substance attaches

itself to the ‘active sites’. This is very often irreversible so prevents

reactant molecules from being adsorbed onto the surface.

For this reason, catalysts have to be regenerated or renewed.

E.g. Lead and its compounds are poisons of transition metal catalysts.

This is why unleaded petrol must be used in cars with catalytic

converters.

Catalysts can also be made ineffective by side-reactions. E.g. Iron used

in the Haber Process eventually rusts.

Active sites


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Enzymes

Enzymes catalyse the chemical reactions which take place in living cells.

Enzymes are complex protein molecules which are very specific- they

usually only speed up one particular reaction and work best at specific

temperatures and pH (optimum).

Examples in nature are:

Amylase – breaks down starch during digestion.

Catalase – breaks down hydrogen peroxide

Many enzymes are used in industry:

Invertase – used in chocolate industry for the hydrolysis of sucrose to

form fructose and maltose.

Zymase - converts glucose into alcohol in the brewing industry.

Protease (and others) – used in biological washing powders to dissolve

natural stains like protein .


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Energy Changes in Chemical Reactions

Chemical reactions involve a change in energy which often results in the

loss or gain of heat energy (exothermic/endothermic reactions)

The heat energy stored in a substance is called its Enthalpy

( H ).

The difference between the enthalpy of the reactants and the enthalpy

of the products in a reaction is the Enthalpy Change

(∆H):

∆H is measured in kJ per mole (kJ mol-1)

Potential Energy Diagrams

We can show the energy changes involved in exothermic and endothermic

reactions by using potential energy diagrams.

A chemical reaction can be regarded as a series of bond breaking and

bond making steps.

Consider the following reaction:


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Exothermic Reactions

Reactions which give out heat energy are called exothermic reactions.

The products have less enthalpy (potential energy) than the reactants

and the temperature of the surroundings increases.

E.g. the combustion of fuels

From this diagram we can work out:

The activation energy (Ea) which is the energy needed to start

the reaction.

The change in enthalpy between the reactants and products

(∆H)

Potential energy

or Enthalpy

kJ mol-1

Reaction pathway

∆H is always negative for exothermic reactions


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Endothermic Reactions

Reactions which absorb heat energy from the surroundings are called

endothermic reactions.

The products have more enthalpy than the reactants and the

temperature of the surroundings decreases.

From this diagram we can work out:

The activation energy (Ea)

The change in enthalpy between the reactants and products (∆H)

∆H is always positive for endothermic reactionsPotential energy

or Enthalpy

kJ mol-1

Reaction pathway


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Activation Energy

The Activation Energy is the ‘energy barrier’ which must be overcome

before the reactants can change into products.

The size of the Activation Energy will control how fast or slow a reaction

is. The higher the Activation Energy (or ‘barrier’) the slower the reaction.

If the activation energy is high, very few molecules will have enough

energy to overcome the energy barrier and the reaction will be slow.

e.g. Combustion of Methane

CH4 + 2O2 CO2 + 2H2O

This is a very exothermic reaction. At room temperature, no reaction

occurs as too few reactant molecules have sufficient energy to react

when they collide. Striking a match provides the molecules with enough

energy to overcome the barrier- it supplies the Activation Energy. Once

started, the energy given out by the reaction keeps it going.

Energy

Products

Reactants

Ea

Large Activation Energy

Slow Reaction


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The Activated Complex

When particles collide with the required Activation Energy (& geometry),

the activated complex is formed.

The activated complex is an unstable intermediate arrangement of

atoms formed as old bonds are breaking and new bonds are forming.

Energy is needed to form the activated complex as bonds in the reactants

may need to be broken, or charged particles brought together.

As the activated complex is very unstable it exists for a very short

period of time. From the peak of the energy barrier the complex can lose

energy to form either the products or the reactants again.

The higher the enthalpy change (∆H), the more unstable the activated

complex.

ReactantsActivated Complex Products


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Catalysts and Activation Energy

A catalyst provides an alternative pathway for the reaction with a lower

activation energy.

N.B The catalyst has no effect on the enthalpy change, ∆H


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PART 2 TRENDS IN THE PERIODIC TABLE

Development of the Periodic Table

• The periodic table was invented by Dimitri Mendeleev (1869).

• He arranged elements in order of increasing atomic

mass, and noted that their properties e.g. Melting point,

boiling point and density were periodic in nature (repeating

patterns existed). .

• Those elements with similar properties were placed below one

another in groups and gaps were left for unknown elements.

• The modern periodic table is based on an elements atomic number,

and this removed a number of the anomalies in the original version.

Trends in Physical Properties of the Elements

Melting points and boiling points

• Melting points and boiling points show periodic properties. This

means that they vary in a regular way or pattern depending on their

position in the Periodic Table.

• Melting points and boiling points depend on the strength of forces

which exist between the particles which make up a substance.

• Going down group 1 the alkali metals M.pt. decrease so there must

be a decrease in the strength of the force of attraction between

the particles.

• Going down group 7 the halogens m.pt. and b.pt. increases so there

must be a increase in the strength of the force of attraction

between the particles

Density

• The density of a substance is its mass per unit volume, usually in

gcm-3. They are found on pg 5 of the data booklet.

• Elements with densities greater than 0.5 gcm-3 are solids and

generally lie to the left-hand side of the Periodic Table.

• Elements with densities less than 0.5 gcm-3 are gases and generally

lie to the right-hand side of the Periodic Table.

In general,


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Covalent Radius - Atomic size

The size of an atom is difficult to measure because atoms do not have a

sharp boundary. However an X-ray technique can be used to measure the

distance between the nuclei of two covalently bonded atoms - this

distance is called the bond length.

The covalent radius of an atom is half the distance between the nuclei of

two of its covalently bonded atoms.

Trends in covalent radii

a. Across a period - covalent radii decreases

Across a period the protons are being added to the nucleus and electrons

are being added to the same shell. The increasing positive charge on the

nucleus pulls the outer electrons more closely and the covalent radius

decreases.

b. Down a group - covalent radii increases

Down a group each member has an

extra shell of electrons so the

covalent radius increases. The

positive charge on the nucleus

increases which tends to pull the

electrons closer but the effect of

adding an extra shell outweighs this.

eg. the bond length in a chlorine

molecule is 198 pm

(pm = picometre: 10-12 metre)

So the covalent radius of a

chlorine atom = 198/2 = 99 pm

These are found on pg 7 of the

data booklet.


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Covalent radius is an example of a Periodic Property as elements in the

same group appear at the same position on the wave.

First Ionisation Energy

This is defined as "the amount of energy required to remove one mole of

electrons from one mole of atoms in the gaseous state”

M (g) M+(g)

+ e 1st ionisation energy

The outermost electron will be the most weakly held and is removed first.

The ionisation energy is an enthalpy change and therefore is measured

per mole. Its units are kJmol-1 (kilojoules per mole). This is always an

endothermic process because energy is required to overcome the

attraction between the electron and the positive nucleus.


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Trends in Ionisation Energy

a. Across a period - ionisation energy increases

Across a period atomic size decreases. This means that the outer

electrons to be removed are closer to the nucleus and are held tightly.

Across a period the size of the positive charge on the nucleus also

increases attracting the electrons more strongly and making them more

difficult to remove.

b. Down a group - ionisation energy decreases

Down a group atomic size increases. This means that the outer electrons

to be removed are further from the nucleus and held less tightly.

However, the size of the nuclear charge is also increasing and you would

expect that the electrons would be strongly attracted to the bigger

nuclear charge. This is not the case. The outer electrons are shielded by

inner energy levels and do not feel the full attraction of the positive

nucleus. This is known as the screening effect. Outer electrons are

easier to remove.


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The second ionisation energy of an element is the energy required to

remove a second mole of electrons.

M+ (g) M2+

(g) + e 2nd ionisation energy

The total energy required to remove 2 moles of electrons is the sum of

the 1st and 2nd ionisation energies.

M (g) M2+

(g) + 2e

The second ionisation energy of an atom is always larger than the first as

it involves removing an electron from a species that is already positively

charged.

Ionisation energies are shown on pg 11 of the data booklet.

Note

1. The large increase from the 1st to the 2nd ionisation energy of lithium. The

1st ionisation energy removes the single outer electron from the lithium atom.

The 2nd ionisation energy requires breaking into a stable electron arrangement -

this requires a lot more energy.

2. With beryllium the largest increase comes between the 2nd and 3rd

ionisation energy. The 3rd ionisation energy requires breaking into a stable shell

of electrons.


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Electronegativity

Electronegativity is a measure of an atom’s attraction for the shared

pair of electrons in a bond.

Which atom would have a greater attraction for the electrons in this

bond and why?

Linus Pauling

Linus Pauling, an American chemist (and winner of two Nobel prizes!) came

up with the concept of electronegativity in 1932 to help explain the

nature of chemical bonds.

Since fluorine is the most electronegative element (has the greatest

attraction for the bonding electrons) he assigned it a value and compared

all other elements to fluorine.

Today we still measure electronegativities of elements using the Pauling

scale.

Values for electronegativity can be found on pg 11 of the data booklet.

Trends in Electronegativity

a. Across a period - electronegativity increases.

This is because the nuclear charge increases, attracting the electrons

more strongly to the nucleus. As a result, the electronegativity increases.

b. Down a group - electronegativity decreases.

Going down the group, the nuclear charge increases but the number of

electron shells also increases. As a result of ‘shielding’ and the increased

distance the outer shell is from the nucleus, electronegativity decreases.


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PART 3 BONDING AND STRUCTURE

Bonds are electrostatic forces (attractions between positive and

negative charges) which hold atoms together.

Atoms form bonds to become more stable - by losing, gaining or sharing

electrons.

The type of bond formed in a substance depends on the elements

involved and their position in the periodic table.

Metallic Bonding

Metallic bonding occurs between the atoms of metal elements.

Metals have little attraction for their outer electrons. These electrons

are free to move so are delocalised.

Electrons can move freely between these partially filled outer shells

creating what is called a ‘sea’ or ‘cloud’ of electrons around positive metal

cores.

The metallic bond is the electrostatic force between positively charged

core and delocalised outer electrons.


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Covalent and Polar Covalent Bonding

Covalent bonding occurs in non-metal elements.

A covalent bond is the electrostatic force of attraction between

positively charged nuclei and negatively charged outer electrons.

In the diatomic element chlorine both atoms have the same

electronegativity so the electrons are shared equally. This is called a pure

or non-polar covalent bond.

In the compound hydrogen iodide the bonded atoms have different

electronegativities. The iodine atom has a bigger attraction for the

shared electrons than the hydrogen atom. As the electrons are attracted

closer to the iodine it becomes slightly negative (δ-) and the hydrogen

atom becomes slightly positive (δ+).

This is called a polar covalent bond.

Other examples are

Ethanol Propanone


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Polar Molecules and Permanent Dipoles

Not all substances with polar covalent bonds will have ‘polar molecules’.

If there is a symmetrical arrangement of polar bonds, the polarity

cancels out over the molecule as a whole and the molecule is non polar.

e.g.

Carbon dioxide Tetrachloromethane

If the bonds are not symmetrical, the molecule has an overall polarity

and is said to have a permanent dipole, i.e. each end has a different

charge.

e.g. Hydrogen Chloride Water

Other examples are; Trichloromethane, ammonia, carbon monoxide


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Covalent Structure

Covalent and polar covalent substances are usually made up of discrete

molecules, but a few have giant covalent network structures.

e.g. Carbon dioxide – discrete molecules

Silicon Dioxide – covalent network structure

(images from BBC Higher Bitesize Chemistry)

Another example of an covalent network compound is silicon carbide.


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The Bonding Continuum

The greater the difference in electronegativity between two elements,

the less likely they are to share electrons, i.e. form covalent bonds.

To judge the type of bonding in any particular compound it is more

important to look at the properties it exhibits rather than simply the

names of the elements involved.

Ionic Bonding

Ionic bonds are formed between metal and non-metal elements with a

large difference in electronegativity.

The non-metal element with the high electronegativity gains the

electrons to form a negative ion:

e.g. Cl + e- Cl-

The element with the low electronegativity loses electrons to form a

positive ion:

e.g. Na Na+ + e-

Both the positive and negative ion will have the same electron

arrangement as a noble gas.

Ionic bonding is the electrostatic force of attraction between

positively and negatively charged ions.

Pure

Covalent

Bond

Polar

Covalent

Bond

Ionic Bond

Increasing ionic character


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Structure of Ionic Compounds

The forces of attraction between the oppositely charged ions results in

the formation of a regular structure called an ionic or crystal lattice.

E.g. Sodium chloride


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Attractions Between Covalent Molecules

There are attractive forces between covalent and polar covalent

molecules which can affect their properties.

These attractions between molecules are called van der Waals or

intermolecular forces (or attractions).

There are 3 types:

1. London Dispersion Forces

2. Dipole-dipole Attractions (permanent dipole-permanent dipole)

3. Hydrogen Bonds are a special type of dipole-dipole attraction which

is particularly strong.

1. London Dispersion Forces

This is the weakest form of intermolecular bonding and it exists between

all atoms and molecules.

Dispersion forces are caused by uneven distributions of electrons.

The atom or molecule gets slightly charged ends known as

a temporary dipole. This charge can then induce an opposite charge in a

neighbouring atom or molecule called an induced dipole. The oppositely

charged ends attract each other creating the intermolecular force. The

relative strength of the force depends on the size of the atoms or

molecules.

Dispersion forces increase with increasing atomic and molecular size.

Uneven distribution

of electrons in Helium

Temporary

Dipole

Induced

Dipole


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2. Permanent Dipole-Permanent Dipole Attractions

A polar molecule is one which has permanently charged ends (permanent

dipole).

Polar-Polar attractions (permanent dipole-permanent dipole) are the

intermolecular force of attraction between the oppositely charged ends

of the polar molecules.

Dipole to dipole attractions are the main forces of attraction between

polar molecules.

Effect of dipole-dipole attractions

Propanone Butane

Formula Mass 58 58

Structure

Main intermolecular permanent dipole- London force

permanent dipole Dispersion

Boiling Point 56oC 0oC

Polar molecules have higher boiling points than non-polar molecules of a

similar mass due to the permanent dipole-permanent dipole interactions.

Permanent dipole-permanent dipole interactions are stronger than London

Dispersion forces.

C

O

CCH

H

H

H

H

H

C C C C

H

H

H

H

H H

H

H

H

H


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3. Hydrogen Bonding

Hydrogen bonds are dipole- dipole interactions found between molecules

which contain highly polar bonds.

They are found in molecules where hydrogen is bonded to very

electronegative atoms like nitrogen, oxygen or flurine (NOF).

Other examples include ammonia, carboxylic acids and alcohols.

Hydrogen bonds are stronger than permanent dipole-permanent dipole

attractions and London dispersion forces but weaker than covalent

bonds.

When Hydrogen bonds are present, the compound will have a much higher

melting point (m.pt) and boiling point (b.pt) than other compounds of

similar molecular size.

Summary of intermolecular attractions

Van der Waals intermolecular forces are much weaker than covalent, ionic

and metallic bonding. The order of strength of the van der Waals forces

is summarised below.

Summary of van der Waals intermolecular forces

London Dispersion Dipole to Dipole Hydrogen

Weakest Strongest


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Bonding and Properties of Elements 1-20

Monatomic Elements - Noble Gases

All consist of single, unbonded atoms.

Only have London Dispersion forces between the atoms.

Properties

Low densities, m.pts and b.pts

Non conductors of electricity as no freely moving charged particles.

B.pts increase as the size of the atom increases.

This happens because more energy is required to break the stronger

London Dispersion forces.

b.p / oC

He

Ne

Ar


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Covalent Molecular Elements (in 1-20)

All consist of discrete molecules of varying size.

Fairly low m.pts, b.pts and densities.

Non-conductors of electricity.

Diatomic elements – H2, N2 , O2 , F2 , Cl2

As the size of the halogen atom increases, so does the

strength of the London dispersion forces. Therefore the boiling point

increase as more energy is required to separate the molecules.

b.p./oC

0

-160

--120

--80

--40

0

40

80

120

160

200

F Cl Br I


Page 41 of 53

Low melting point solids.

Phosphorus – P4

Sulphur - S8

Higher m.pt. because there are stronger London Dispersion forces

between larger molecules.

Fullerenes (Carbon)

Buckminster fullerene C60 (Bucky Balls) discovered in the 1980’s

Due to the large molecules , fullerenes have stronger dispersion forces

between their molecules than smaller

molecules.

NB – they are molecules not covalent networks

Nanotubes


Page 42 of 53

Covalent Network Elements (in 1-20)

Giant network structures containing millions of

atoms.

E.g. Carbon exists in 2 main forms…

Diamond

4 bonds per carbon atom – tetrahedral

structure

Non-conductor of electricity as no free

electrons (delocalised).

Hardest natural substance as many strong

bonds to break so used for drills, cutting

tools, etc.

Graphite

3 bonds per carbon atom – layered structure

with London dispersion forces between the

layers

Conductor of electricity due to one

delocalised electron per atom able to move

between the layers – used in electrodes.

Very soft – the layers break away easily

due to weak dispersion forces so good as

a lubricant and for drawing (pencils).


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Metallic elements (Revision of Nat 5)

All have metallic lattice structure.

The metallic bond is the attraction between

the positive metal cores and the outer

delocalises electrons.

Metallic bonds are generally strong. The strength of the bond will

depend on the number of outer electrons that each atom contributes to

the delocalised pool. The greater the number of outer electons, the

stronger the metallic bond.

They conduct electricity when solid or liquid due to free moving

delocalised outer electrons.

Positive core

Delocalised electrons

http://www.google.co.uk/url?sa=i&source=images&cd=&cad=rja&uact=8&docid=RUkvxDAIrxn0nM&tbnid=VKyvh3y1-Ld-KM:&ved=0CAgQjRw&url=http://www.bbc.co.uk/schools/gcsebitesize/science/add_ocr_gateway/periodic_table/metalsrev2.shtml&ei=itB5U9THOai70QWwxYCYBA&psig=AFQjCNGq_hMIdxunis1e7dad0A8MA3g2ig&ust=1400578571096609


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The first twenty elements

The first twenty elements of the Periodic Table can be grouped according

to the type of bonding and structure.

Covalent molecular

A covalent molecular structure consists of discrete molecules held

together by London dispersion forces of attraction. Some elements

normally exist as solids, others exist as diatomic gases.

Covalent network

A covalent network structure consists of a giant lattice of covalently

bonded atoms.

Metallic

A metallic structure consists of a giant lattice of positively charged ions

and delocalised outer electrons.

Monatomic A monatomic structure consists of discrete (separate) atoms

held together my van der Waals' forces of attraction.

Complete the following table by including the appropriate type of bonding

and structure

A

B

C(i)

C(ii)

D


Page 45 of 53


Page 46 of 53

Bonding and Properties of Compounds

Compounds can be split into 3 main groups, depending on their bonding,

structure and properties:

1. Ionic Lattice Structures

2. Covalent Network Structures

3. Covalent Molecular Structures

1. Ionic Lattice Structures

All ionic compounds are solids at room

temp so have high melting and boiling

points. This is because the ionic bonds

holding the lattice together are strong

and a lot of energy is required to

break them.

The stronger the ionic bond the higher

the melting point.

Ionic compounds conduct electricity

when dissolved in water or when molten as the ions are free to move.

Electrolysis of an ionic solution or melt causes a chemical change at the

electrodes.

They do not conduct when solid as the ions are ‘locked in the lattice and

cannot move to carry the current.

2. Covalent Network Structures

Covalent networks have very high melting and boiling points as many

strong covalent bonds need to be broken in order to change state. They

can also be very hard.

Silicon Carbide (SiC) – carborundum, has a similar structure to diamond.

It has a high melting point (2700oC). and it is used

as an abrasive.

CovalentBondTetrahedral

shape


Page 47 of 53

The 4 carbon atoms are available to bond with

another 4 silicon atoms resulting in a

covalentnetwork.

Covalent network structures are usually non-

conductors of electricity as they have no free

moving charged particles.

2. Covalent Molecular Structures

Usually have low melting and boiling points as there is little attraction

between their molecules.

E.g. Carbon dioxide CO2: m.pt -57oC (non-polar)

Compounds with polar molecules may have slightly higher m.pts and b.pts

than non-polar molecules due to permanent dipole-permanent dipole

attractions

e.g. Iodine chloride Bromine

I - Cl Br – Br

b.pt 97oC b.pt 59oC

Effects of Hydrogen bonding on properties.

1. Melting and boiling points

When hydrogen bonds are present, the compounds will have a much

higher m.pt and b.pt than other compounds of similar molecular size as

more energy is required to separate the molecules.

Water has a much higher b.pt than similar compounds containing hydrogen

= Carbon

= Silicon


Page 48 of 53

Hydrogen bonding explains why water, HF and NH3 have a b.pt higher than

expected.

Similarly HF b.p. 19 oC Whereas: HBr –68 oC and HI –35 oC

2. Viscosity

Viscosity is how thick a liquid is. As molecules get bigger the viscosity

increases.

Hydrogen bonding also affects the viscosity. The more hydrogen bonding

between molecules (ie the more OH groups present) the more viscous the

liquid.

Substance ethanol water glycerol

Molecular mass 46 18 92

Structural Formula

No of –OH groups 1 2 3


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3.Miscibility

Miscible liquids mix thoroughly. Ethanol and water would be described as

miscible but water and oil are immiscible as the oil forms a visible layer

on water.

Hydrogen bonding aids miscibility and ethanol and water both contain

hydrogen bonds so mix easily.

4.Density

Why do pipes burst when water freezes and why does ice float on

water?

As matter is cooled, it normally contracts and becomes more dense.

However, as water freezes it expands (at about 4oC) because the strong

hydrogen bonds between the molecules force them into an open lattice

structure.

This makes the solid ice less dense (takes

more space) than the liquid so ice floats

on water and pipes burst when water

freezes.


Page 50 of 53

Bonding, Solubility and Solutions

Ionic lattices and polar covalent molecular compounds tend to be:

Soluble in water and other polar solvents, due to the attraction

between the opposite charges.

Insoluble in non-polar solvents, as there is no attraction between

the ions and the solvent molecules.

Non-polar covalent molecular substances tend to be:

Soluble in non-polar solvents like carbon tetrachloride or hexane.

Insoluble in water and other polar solvents as there are no

charged ends to be attracted.

-

++

-

++- ++ -+ -

Water molecule

Ionic lattice

Hydratedions

--

+

+ -+

+

-

+ +

+-+

+

+ -

+

-

+ +

+ve ions attracted to –ve ends of water molecule

-ve ions attracted to +ve ends of water molecule


Page 51 of 53

Rate of Reactions - Glossary

Word Meaning

Activation energy (Ea)

The minimum amount of energy needed for a reaction to begin.

Catalyst A chemical which speeds up a chemical

reaction without being used up itself and which can be removed

chemically unchanged at the end of the reaction.

Catalytic converter A catalyst found in the exhaust of cars. It changes harmful

gases into less harmful gases.

It is usually made of platinum.

Chemical reaction An interaction between substances (chemicals) in which their

atoms re-arrange to form new substances.

Concentration The amount of particles in a given volume.

Enzyme A biological catalyst (ie. a catalyst found in living things).

Products The substances (chemicals) at the end of a chemical reaction.

Rate of reaction How quickly a reactant is used up OR how quickly a product is

created.

Surface area Total area of a substance which is exposed to the surroundings.

Heterogeneous The reactants are in a different state from the catalyst (the

catalyst is generally a solid).

Homogeneous The reactants and the catalyst are in the same state

Kinetic Energy The movement energy of particles.

Potential Energy The stored energy in reactants or products.

Activated Complex

a very unstable intermediate arrangement of atoms where the old

bonds are being broken and the new bonds are forming

Exothermic Giving out heat to the surroundings (feels hot). ∆H= -kJmol-1

Endothermic Taking heat in from the surroundings (feels cool). ∆H= +kJmol-1


Page 52 of 53

Trends in the Periodic Table - Glossary

Word Meaning

Covalent radius Half the distance between the nuclei of two bonded atoms of an

element.

Density The density of a substance is its mass per unit volume, usually in

gcm-3.

Electronegativity A measure of the attraction that an atom involved in a covalent

bond has for the electrons of the bond.

First ionisation energy The amount of energy required to remove one mole of electrons

from one mole of atoms in the gaseous state.

Periodicity The occurrence of patterns in the Periodic Table.

Screening (shielding)

effect

When inner electrons shield an outer electron from the

attractive effect of the nucleus and less energy is needed to

remove the outer electron as a result.

Bonding and Structure - Glossary

Word Meaning

Chemical bonding is the term used to describe the mechanism by which atoms are

held together.

Chemical structure describes the way in which atoms, ions or molecules are arranged.

Covalent bond a covalent bond is formed when two non-metal atoms share a pair

of electrons .

Covalent radius is half the distance between the nuclei of two bonded atoms of

an element Delocalised Delocalised electrons, in metallic bonding, are free from

attachment to any one metal ion and are shared amongst the

entire structure. Dipole an atom or molecule in which a concentration of positive charges

is separated from a concentration of negative charge.

Fullerenes are molecules of pure carbon constructed from 5- and 6-

membered rings combined into hollow structures. The most stable

contains 60 carbon atoms in a shape resembling a football.


Page 53 of 53

Word Meaning

Hydrogen bonds are electrostatic forces of attraction between molecules

containing a hydrogen atom bonded to an atom of a strongly

electronegative element such as fluorine, oxygen or nitrogen, and

a highly electronegative atom on a neighbouring molecule. Intermolecular forces Forces of attraction between molecules (or atoms)

Ionisation energy The first ionisation energy is the energy required to remove one

mole of electrons from one mole of gaseous atoms (i.e. one

electron from each atom). The second and subsequent ionisation

energies refer to the energies required to remove further moles

of electrons. Isoelectronic means having the same arrangement of electrons. For example,

the noble gas neon, a sodium ion (Na+) and a magnesium ion (Mg2+)

are isoelectronic. Lattice A lattice is a regular 3D arrangement of particles in space. The

term is applied to metal ions in a solid, and to positive and

negative ions in an ionic solid. London Dispersion

Forces are the intermolecular forces of attraction which result from the

electrostatic attraction between temporary dipoles and induced

dipoles caused by movement of electrons in atoms and molecules. Miscible fluids which mix with or dissolve in each other in all proportions.

Polar covalent bond a covalent bond between atoms of different electronegativity,

which results in an uneven distribution of electrons and a partial

charge along the bond. Non-polar (pure)

covalent bond a covalent bond between atoms of the same electronegativity,

which results in an even distribution of electrons .

van der Waals forces Is the general name given to all intermolecular attractions

including London dispersion, permanent dipole to permanent dipole

to dipole attractions and hydrogen bonding. Viscosity The thickness of a liquid.

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Unit 1 Chemical Changes & Structure - Duncanrig Secondary School · 2017. 6. 27. · Duncanrig Secondary School cfe Higher Chemistry Chemical Changes and Structures Page 3 of 53 Learning