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CfE Higher (National 6) - Unit 1 Notes Unit 1a - Controlling the rate Revision of previous knowledge (National 5) Before a chemical reaction can take place a collision must take place between reactant particles:

A + B --> C

The more successful collisions the faster the reaction. (successful collisions will be discussed in more detail later in these notes). At National 5 you learned that the following factors could influence the speed of a chemical reaction: CONCENTRATION - increasing the number of reactant particles increases the speed of the reaction. SURFACE AREA - as only particles on the surface of a solid react, powders, with their large surface area, react faster than lumps. TEMPERATURE - the higher the temperature the faster the reaction. CATALYSTS - offer an alternative route (a short cut) that requires less energy. We look at the factors of TEMPERTURE and CATALYSTS in more detail later in this topic. All chemical reactions involve: Reactant ------> Product, i.e. R --------> P Reactions can be followed by noting changes in concentration of reactant (R) or product (P) with time, e.g. if the product is a gas its volume could be measured with a gas syringe, or, a balance could be used to record the loss in mass of the reactant vessel. The experimental results are often presented as line graphs:

The rate of reaction is defined as the change in concentration of reactant or product in unit time. The rate equation: Rate = Change in concentration of R or P = mol l-1 = mol l-1 s-1

Time for change s If a reaction was followed by measuring the volume of gas released, the rate equation would be: Rate = Change in volume = cm3 = cm3s-1 Time for change s

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If a reaction was followed by measuring the loss in mass, the rate equation would be: Rate = Change in mass = g = gs-1 Time for change s

Collision Theory Not all collisions result in a chemical reaction. There are TWO requirements for a successful collision: 1. The particles must collide at the correct angle - COLLISION GEOMETRY. Particles in a reaction will only collide to form the product if they are correctly orientated to each other. Consider the reaction: NO (g) + O3 (g) ----> NO2 (g) + O2 (g)

If the particles collide at the wrong angle the molecules simply bounce off one another. If the collision geometry is correct a successful collision will occur and the products will form. 2. The particles need a certain minimum energy - ACTIVATION ENERGY (EA) A reaction can only take place if the particles are moving fast enough (sufficient kinetic energy) to break the bonds between atoms/molecules. This need for a minimum amount of energy is a ‘barrier’ that prevents a successful collision. The energy required for a successful collision is called the activation energy (EA). The energy of activation can be achieved by:

• Heat - (most common) - raises the average kinetic energy of the reactants.

• Electric spark - igniting natural gas or petrol in a car’s engine.

• Light - photosynthesis.

Relative rate As a rapid reaction is over in a short time, the rate of a reaction is inversely proportional to the time of the reaction. Rate = 1 = 1 = s-1

t s

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Activation energy and temperature

A graph of reaction rate against temperature looks like: From the graph you can see that the reaction is very sensitive to temperature. A small rise in temperature can result in a large increase in reaction rate, e.g. in some reactions the reaction rate doubles for every 10⁰C temperature rise. As such, the effect of temperature on reaction rate can not simply be explained in terms of an increase in the number of collisions with a rise in temperature. Reactant particles have mass and are moving as a result they have kinetic energy. This energy varies from particle to particle, i.e. some are moving faster than others. The distribution of kinetic energies among the particles at a given temperature (20⁰C) can be represented by an ENERGY DISTRIBUTION DIAGRAM like the one shown below: Although the kinetic energies of individual particles are always changing, e.g. with collisions, the overall picture stays the same. Only a very small fraction of the particles have very low energies and similarly only a very small fraction have high energies. The majority have an energy equal to or closely to the average kinetic energy. The average kinetic energy of the particles is directly proportional to the temperature of the substance. Therefore temperature is widely defined as: “a measure of the average kinetic energy of the particles” In the energy distribution diagram the line marked EA represents the activation energy of a reaction. The shaded area represents those particles which have the necessary activation energy at 20⁰C for a successful collision. If the temperature of the gas is increased to 90⁰C then the fraction of the particles with very low energies decreases and the fraction of particles with very high energies increases. This is shown in the diagram below along with the arrangement of particles at the lower temperature (20⁰C): The shaded area represents those particles that have the necessary activation energy at 90⁰C. Clearly, as the temperature increases more reactant particles have an energy equal to or greater than the activation energy for the reaction.

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Revision of previous knowledge (National 5) When a chemical reaction takes place energy is required to start changing the reactants into products and energy is released as the products start to form. An EXOTHERMIC REACTION is a reaction which releases heat energy to the surroundings causing the temperature of the surroundings to rise.

EXO --> EXIT THERM --> HEAT As a result exothermic reactions are often thought of as HOT reactions.

An ENDOTHERMIC REACTION is a reaction which absorbs (takes in) heat energy from the surroundings causing the temperature of the surroundings to fall.

ENDO --> ENTER THERM --> HEAT

As a result endothermic reactions are often thought of as COLD reactions.

Enthalpy Every substance contains stored energy known as chemical energy. The energy content if a substance is known as ENTHALPY (symbol H). Enthalpy varies from substance to substance and so, during a chemical reaction, there is a change in enthalpy

(symbol H), detected as a heat change, when reactants (R) go to products (P).

The enthalpy change for a reaction can be calculated: H = HP - HR (Where HP and HR are the enthalpies of the products and reactants respectively) If the enthalpy of the reactants is greater than the enthalpy of the products then energy will be released (given out)

during the reaction. This is called an exothermic reaction. H will have a negative value and will result in an increase in temperature during the reaction. This is represented by the diagram shown below. If the enthalpy of the products is greater than the enthalpy of the reactants then energy will be absorbed (taken in)

during the reaction. This is called an endothermic reaction. H will have a positive value and will result in a decrease in temperature during the reaction. This is represented by the diagram shown below.

The diagrams above, however, do not show the complete picture. As the reactant molecules approach each other they must possess sufficient kinetic energy to overcome the repulsive forces between electron clouds and to weaken or break the bonds within the molecules. (Kinetic energy --> Potential energy). This minimum amount of kinetic energy is known as the activation energy (EA). At the point of impact, the potential energy (P.E), is at a maximum. As the products move away, the PE falls again (PE --> KE). To represent this chemists use potential energy diagrams.

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Potential energy diagrams Potential energy diagrams give useful information about the energy profile of a reaction. When drawn accurately they can be used to calculate the enthalpy change and/or the activation energy of a reaction. These diagrams are known as potential energy diagrams as the enthalpy change for a given reaction will be the same as the difference in PE.

At the top of the PE barrier (the maximum PE) an activated complex forms. As a result of the high PE, this is an un-stable arrangement of atoms and is a short lived intermediate stage between reactants and products. To reach this stage, the reactants must possess a minimum amount of kinetic energy initially. The energy required to form the activated complex is known as the activation energy (EA) for the reaction and is the energy difference between reactants and the top of the barrier (if the reaction is reversed, the EA for that reaction will be the difference between the products and the top of the barrier). Once the activated complex forms it can either re-form the reactants by releasing energy equivalent to the activation energy of form products by releasing more energy (an exothermic reaction) or less (an endothermic reaction) energy that the activation energy. Lets consider the decomposition of hydrogen iodide into its elements: 2HI (g) H2(g) + I2 (g) As the hydrogen iodide molecules collide, the H-I bonds are weakened and partial bonds are set up between the H atoms and between the I atoms. (Remember the importance of correct collision geometry?) Showing this on a PE diagram we get the following:

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Determining activation energies and enthalpy changes from PE diagrams Example 1: Consider the potential energy diagram: The activation energy for the forward reaction can be determined by measuring the gap between the line for the reactants marked X to the top of the curve. From the shape of the diagram we can see that the forward reaction (X --> Y) is endothermic. The reactants have less energy than the products. To determine the enthalpy change for the reaction we measure the gap between the line marked X and the line marked Y.

To determine the activation energy for the reverse reaction we simply subtract the enthalpy change H from the activation energy for the forward reaction EA(FORWARD). Because the activation energy is the minimum amount of energy required to start the reaction this value will always be positive. EA(FORWARD) = 103kJ

H = + 57kJ EA(REVERSE) = 103 - 57 = 46kJ Example 2: Consider the potential energy diagram: The activation energy for the forward reaction can be determined by measuring the gap between the line for the reactants marked A to the top of the curve. The activation energy for the reverse reaction can be determined by measuring the gap between the line for the reactants marked B to the top of the curve. From the shape of the diagram we can see that the forward reaction (A --> B) is exothermic. The reactants have more energy than the products. To determine the enthalpy change for the forward reaction we simply subtract the activation energy for the reverse from the activation energy for the forward reaction EA(FORWARD) = 142kJ EA(REVERSE) = 201kJ

H = 142 - 201 = -59kJ Example 3: Consider the potential energy diagram: The activation energy for the forward reaction can be determined by measuring the gap between the line for the reactants marked A+B to the top of the curve. EA(FORWARD) = 60kJ The activation energy for the reverse reaction can be determined by measuring the gap between the line for the reactants marked C+D to the top of the curve. EA(REVERSE) = 80kJ

To determine the enthalpy change for the forward reaction we

simply subtract the potential energy of the products from the potential energy of the reactants. H = 20 - 40 = -20kJ The negative enthalpy change means the reaction is exothermic.

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Revision of previous knowledge (National 5) At national 5 we learned about a group of chemicals that had a positive effect on the speed of a chemical reaction. These chemicals we called CATALYSTS. A catalyst is any chemical that:

• Increases the rate of a chemical reaction

• Can be recovered chemically unchanged at the end of the chemical reaction

Catalysts A catalyst speeds up a chemical reaction by providing a different pathway with a lower activation energy. As a result more molecules (particles) have sufficient kinetic energy (KE) to overcome the activation energy. The above energy distribution diagram shows how the Ea for the catalysed reaction is significantly lower than that of the uncatalysed reaction. Whilst the reactants at the start and the products at the end are the same, during the reaction the catalyst (Z) alters the reaction pathway and produces different activated complexes.

The uncatalysed reaction pathway

The catalysed reaction pathway As a result of the reactants and products being the same for both the catalysed and uncatalysed reaction the

enthalpy change (H) will remain the same. This can be seen clearly on a potential energy diagram shown below:

Note: A catalyst: 1. Lowers the activation energy for a reaction 2. Leaves the enthalpy change, H, unchanged

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Catalysts are often divided into two groups:

HOMOGENEOUS and HETEROGENEOUS A homogeneous catalyst is in the same state as one of the reactants. You have already met and will meet again this year esterification reactions. Esters are formed from liquid alcohols and carboxylic acids. The catalyst is the liquid concentrated sulphuric acid. As the catalyst is in the same state as the reactants it is a homogeneous catalyst. A heterogeneous catalyst is in a different state from the reactants, e.g. reactions involving gases or liquids being catalysed by solids. The catalyst is often a surface catalyst as the reaction takes place on the catalysts surface. The reactive parts on the catalysts surface are called ACTIVE SITES.

1. (ADSORPTION) Molecules of one or both reactants form bonds with the catalyst. This weakens the bonds within the molecules.

2. (REACTION) The molecules react on the catalyst surface. The angle of collision is more likely to be favourable since one of the molecules is fixed.

3. (DESORPTION) The product molecules leave the catalyst and the vacant active site can be occupied by anoth-er reactant molecule.

A heterogeneous catalyst speeds up the reaction by: 1. Attracting the reactant particles to its surface, thus increasing the concentration of reactants. 2. The reactant particles are adsorbed on the active site forming bonds with the surface atoms. As these bonds

form, the bonds within the particles are weakened. Less energy is now required to break the bonds so the activation energy has been lowered.

3. As the adsorbed particles are held on the active site in a fixed position the collision geometry is improved. If impurities are present in the reactants they may be adsorped onto the active site of the catalyst, e.g. lead poisons the platinum catalytic converter in a cars exhaust system. If the reactant particles cannot be adsorped onto the catalysts active site then the catalyst will no longer function. It is important during industrial processes that reactants be pure and free of any foreign material. For example, in the Haber process, the iron catalyst is susceptible to poisoning by carbon monoxide. The vanadium (V) oxide (V2O5) catalyst in the contact process (used to make sulfuric acid) is readily attacked by arsenic compounds. Industrial catalysts you will have met at national 5 include: Iron - The Haber process - the manufacture of ammonia from nitrogen and hydrogen Platinum - The Ostwald process - the manufacture of nitric acid from ammonia Aluminium oxide - Cracking hydrocarbons - breaks long chain hydrocarbons into shorter more useful ones Concentrated sulfuric acid - Esterification - making esters from alcohols and carboxylic acids

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Unit 1b- Periodicity Revision of previous knowledge (National 5) The Periodic Table The modern periodic table is based on the work of the Russian chemist Dimitri Mendeleev. Mendeleev put the element in order of increasing atomic mass but left gaps in his table for elements still to be discovered. Mendeleev used his table to predict the properties of the elements still to be discovered. His predictions were very accurate. The modern Periodic Table is based on increasing atomic numbers.

Elements are arranged in rows and columns in the Periodic Table. A row is called a period and a column is called a group.

Both periods and groups are numbered in the periodic table and certain groups are given unique names. For instance group 1 are called the alkali metals. All the elements in this group are highly reactive as a result of having a single outer electron. Mendeleev deliberately placed elements with similar chemical properties in groups.

The group 7 non-metal elements are called the halogens. Like the alkali metals these too are very reactive elements this time due to being 1 electron short of a full outer energy level (the halogens all have 7 outer electrons).

The noble gases (group 0) are an unreactive group of non-metals as they have full outer energy levels.

Bonding in elements - Metals

When metal atoms combine a metallic bond forms. The bond is formed by the nuclei of neighbouring atoms attracting the delocalised elec- trons.

Metallic bonding is describe as positive ions in a “sea of delocalised electrons”

Metallic bonds are strong and as a result the majority of metals are solids at room temperature. The delocalised or freely moving electrons mean that all metals are conductors of electricity in both the solid and liquid state.

Delocalised outer energy level electrons

Nuclei and inner energy level electrons, i.e.

positively charged ions

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Bonding in elements - Non-metals

When non-metal atoms combine, the stable noble gas electron arrangement is achieved by atoms sharing their bonding electrons. Non-metals are held together by a covalent bond.

The shared electrons (negative) are attracted by both nuclei (positive) so forming a covalent bond. This prevents the electrons from moving freely and as a result all covalently bonded elements are insulators of electricity.

A lot of energy is required to overcome this force of attraction and so we can conclude that covalent bonds are also strong. A covalent bonds occur within molecules, we call them intramolecular bonds.

There are two types of covalent structure: MOLECULAR (DISCRETE) COVALENT and NETWORK COVALENT.

A covalent molecular element consists of independent molecules, weakly held together by intermolecular forces known as London dispersion forces.

Covalent molecular elements have relatively low melting and boiling points because when a covalent molecular element is melted it is the weak intermolecular forces not the strong intramolecular, covalent bonds being broken.

Some covalent elements have a network covalent structure. Unlike covalent molecular all the atoms in a covalent network element are held together by strong covalent bonds and as a result have very high melting and boiling points.

The noble gases (group 0 or 8) despite being non-metals do not form covalent bonds. Instead the noble gases exist as monatomic gases held together by weak intermolecular forces.

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Structures and properties of the first 20 elements Metals (Li, Be, Na, Mg, Al, K, Ca) All metals conduct electricity in the solid and molten states as the delocalised electrons are able to move freely. The melting and boiling points are high indicating that metallic bonds are strong. The melting point of the metal elements increase as you move right across a period in the Periodic Table. This is because the number of delocalised electrons increase as you move right along a period, i.e. Na - 2, 8, 1; Mg - 2, 8, 2 and Al- 2, 8, 3. This increases the strength of the metallic bond and raises the melting and boiling point. On going down a group the melting and boiling points of the metal elements decrease. This indicates that the strength of the metallic bond also decreases descending a group. This is because despite each atom having the same number of outer electrons the number of energy levels increases. This means the outer electrons are further from the nucleus and as a result are held in place less tightly, decreasing the strength of the metallic bond, i.e. Li - 2, 1; Na - 2, 8, 1 and K - 2, 8, 8, 1. Non-metals The properties of the non-metals elements depend on their structures: covalent molecular, covalent network and monatomic. The trend in melting and boiling points going across the Periodic Table confirms your knowledge of bonding and structure from National 5. Covalent molecular (H2, N2, O2, F2, Cl2, P4, S8, Fullerenes) All the elements that exist as covalent molecular elements are electrical insulators as there are no freely moving delocalised electrons. They all have low melting and boiling points as the bonds broken are the very weak intermolecular London Dispersion Forces (You will learn more about the formation of London Dispersion Forces in unit 1c). The strength of the intermolecular force depends on the size and mass of the molecule. As you go down a group of covalent molecular elements (like the Halogens) in the Periodic Table, the melting and boiling point increases as the size of the molecule also rises. The melting point and boiling point decreases from nitrogen to fluorine (as you move right across a period) as the size of the molecule decreases. Some elements have a higher melting point than expected, e.g. sulfur and phosphorus. These molecules are not diatomic - instead they exist as larger molecules P4 and S8 (sometimes called a cluster of atoms) with stronger intermolecular forces. As was seen with the diatomic elements, the bigger the molecule the higher the MP/BP. The discrete molecular solids, i.e. sulfur and phosphorus are both soft indicating the weak nature of the intermolecular force. Carbon can exist in many forms (polymorphs). You have met carbon as 2 distinct network forms but carbon also ex-ists covalent molecular forms known as Fullerenes. In Fullerenes the carbon atoms join together to make large mole-cules, e.g. C60, C70 etc. The first Fullerene is called Buckminsterfullerene. All Fullerenes are examples of covalent molecular elements and as such consist of covalently bonded molecules held together by intermolecular London dispersion forces. Fullerenes are the subject of much scientific research in the hope of opening new areas in chemistry such as the development of nanotechnology.

C60

S8 P4

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Covalent network (B, C, Si) The covalent network elements boron, carbon and silicon have different properties from the covalent molecular ele-ments. The structures are held together by very strong covalent bonds. Carbon (diamond), boron and silicon have the following structure: Covalent networks like Carbon diamond have no free moving electrons so they do not conduct electricity. The melting points and boiling points are very high, C - mp = 3925K and Si - mp = 1683K, compared with molecular cova-lent elements, reflecting the strength of the covalent bond. Boron, carbon and silicon are very hard. Carbon (graphite) has the following structure: In graphite each carbon atom is covalently bonded to only three other carbon atoms. The layers of carbon atoms are held together by weak London dispersion forces. As such the layers can slide over each other making graphite soft and slippery. With only 3 of the 4 outer electrons involved in bonding graphite is a good electrical conductor as it has free moving, delocalised electrons, between the layers. Like diamond the melting point is very high as it still contains many millions of atoms covalently bonded together. Monatomic (He, Ne, Ar) To the far right hand side of the Periodic Table are a group of elements called the NOBLE GASES. These non-metallic elements are characteristically unreactive and are thought of as having atoms that do not bond to each other at all. As a result they are also known as the INERT or MONATOMIC ELEMENTS. However, if a noble gas like argon is cooled, as you would expect from a gas it will; if cooled enough; condense and change state into a liquid. This liquid could then be changed into a solid, if the freezing point of argon was exceeded. As a liquid or solid a weak attraction must be holding the argon atoms together. This weak attraction is known as a LONDON DISPERSION FORCE. London dispersion forces are weak attractions that exist between atoms or molecules. It is this weak attraction that explains the very low melting point of this group of unreactive non-metals.

Boron Silicon Diamond

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Trends in the Periodic Table

Atomic/Covalent radius The atomic/covalent radius of an element is half the distance between the nuclei of two of its bonded atoms. On crossing the Periodic Table (left to right) the atomic radius/covalent radius decreases. This is due to the increased nuclear attraction for the outer electrons. On going down a group in the Periodic Table the size of the atoms increases. This is because on moving from one element to the next, the number of occupied energy levels increases. The increased nuclear charge is cancelled out by the screening or shielding effect of the inner energy levels. Ionisation energy (I.E.) The reactivity of an element depends on the electrons in their atoms and how they are arranged. An important measure of an atom is the amount of energy required to pull electrons off an atom and the resulting formation of ions. Since removing electrons causes an imbalance in the protons and neutrons the resulting particle will carry a charge and is therefore described as an ION. The amount of energy required to remove an electron from an atom is therefore called the IONISATION ENERGY (I.E.) or the ENTHALPY OF IONISATION. The first ionisation energy is the amount of energy required to remove one mole of electrons from one mole of gaseous atoms, e.g.

K (g) --> K+(g) + e- H= 425kJmol-1

It is an endothermic reaction as energy is required to break the bond of attraction between the nucleus and the out-er electron. The closer an electron is to the nucleus (i.e. the smaller the atom) the higher the ionisation energy. On crossing the periodic table (left to right) the size of the atoms gets smaller due to the increased nuclear charge. Both effects make it more difficult to remove an electrons so the ionisation energy increases. On going down a group the ionisation energy decreases. As the atoms get larger the outer electron moves further from the nucleus - less attraction, so lower ionisation energy. The effect of the increased nuclear charge is cancelled out by the screening or shielding effect of the inner energy levels. The formation of ions often involves the loss of more than one electron, e.g. Mg2+ or Al3+, etc, so second, third and forth ionisation energies exist for many elements.

Example 1: Na (g) --> Na+(g) + e- H= 502kJmol-1 , first ionisation energy

Na+ (g) --> Na2+(g) + e- H= 4650kJmol-1 , second ionisation energy

The second ionisation energy is far greater than the first. To remove the second electron involves breaking into a full, stable energy level (like that of a noble gas), Na (2, 8, 1) --> Na+ (2, 8) --> Na2+ (2, 7).

This requires so much energy it is unlikely to occur.

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Example 2: Mg (g) --> Mg+(g) + e- H= 744kJmol-1 , first ionisation energy

Mg+ (g) --> Mg2+(g) + e- H= 1460kJmol-1 , second ionisation energy

Mg2+ (g) --> Mg3+(g) + e- H= 7750kJmol-1 , third ionisation energy

The third ionisation energy is far greater than the first. To remove the third electron from the Mg2+ ion involves breaking into a full, stable energy level (like that of a noble gas), Mg2+ (2, 8). The second ionisation energy is greater than the first as an electron is being removed from a positive ion, thus the effect of the net nuclear charge on the electron is greater. The energy required to form the Mg2+ ion is the sum of the first and second ionisation energies:

Mg (g) --> Mg2+(g) + 2e- H= 744 + 1460 = 2204kJmol-1

Additional points on Ionisation energy (I.E.) 1. The first ionisation energy of magnesium is greater than that of sodium magnesium is smaller due to its larger

nuclear charge. Therefore there is a greater attraction for the outer electron. 2. Ionisation energy is an indicator of the reactivity of metals. Many metals react by losing their outer electron.

Caesium has the lowest ionisation energy stated in the data book, so is therefore the most reactive. Electronegativity The attraction of an atom for electrons depends on its size and nuclear charge. Electronegativity is a measure of the attraction of an element for the shared electrons in a covalent bond. Small atoms with a high net nuclear charge (excluding the shielding effect) will attract electrons more than larger atoms. In general the electronegativity increases as you move right along a period since the nuclear charge increases in the same direction, and the electronegativity decreases down a group of the Periodic Table since the atomic size in-crease down the group.

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Unit 1c- Bonding and structure Revision of national 5 Metal alloys When metals combine with other metals (and sometimes non-metals) they form an alloy, e.g. steel is an alloy of iron and carbon. An alloy is not a compound but a mixture of elements. The properties of alloys are based on the metallic bond so alloys have similar and often improved properties of the parent metal, e.g. stainless steel does not rust despite containing iron. Metal and non-metal compounds - Ionic bonding Ionic bonding is the electrostatic attraction between the positive ions of one element and the negative ions of another. Ionic compounds have high melting and boiling points indicating a strong bond between oppositely charged ions. They exist as ionic lattices and are all solids at room temperature. Ionic compounds do not conduct electricity in the solid state as the ions are not free to move. However in the molten state or in solution the ions are free to move, so they will conduct electricity in a liquid or aqueous state. Non-metal compounds - Covalent bonding As was seen with the covalently bonded elements, 2 structures exist within the covalent compounds. Some covalent compounds exist as covalent networks. Examples include silicon dioxide (sand) and silicon carbide. Both substances have very high melting and boiling points and are extremely hard. This is because when covalent network compounds are melted the strong covalent bonds between atoms need to be broken. Silicon dioxide is used to coat sand paper for smoothening the surface of wood, whilst silicon carbide is used widely for the discs of cutting tools like angle grinders. The other structure that exists within covalent compounds is covalent molecular. Unlike covalent networks, covalent molecular compounds tend to be liquids or gases at room temperature as a result of their relatively low melting and boiling points. This is because when a covalent molecular compound is melted or boiled it is not the strong covalent bonds being broken, but instead weak intermolecular forces.

Ionic bonding Ionic bonding is the electrostatic attraction between the positive ions of one element and the negative ions of another. The bigger the difference in electronegativity between the two bonding elements the less likely they are to share electrons. This means it is unlikely that compounds consisting of elements with large differences in electronegativity will form covalent bonds. Instead the element with the greatest pull for shared electrons (electronegativity) will completely remove both bonding electrons and create a full outer energy level for itself. This is now a negatively charged ion. The element with the lower electronegativity will lose its bonding electron forming a positive ion.

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Non-Polar (Pure) and Polar covalent bonding Covalent molecular compounds consist of non-polar (pure) or polar molecules. Whether a molecule is polar or pure depends on:

• The difference in electronegativities of the 2 bonding elements

• The shape/symmetry of the molecule Polar covalent bonds Many covalent compounds consist of elements with different electronegativities. Earlier in this unit we learned that the electronegativty of an element is the amount of attraction it has for the shared electrons within a covalent bond. When a compound contains 2 or more elements with significantly different electronegativities the bonding electrons are not shared equally. The atom with the greater electronegativity pulls the bonding electrons closer to itself resulting in those atoms becoming slightly negative and the atoms with the lower electronegativity becoming slightly positive. The symbols δ+ and δ- means slightly positive and slightly negative respectively. This arrangement where the electrons are shared unevenly is known as a polar covalent bond. Water and hydrogen fluoride are 2 common examples of molecules containing polar covalent bonds: A molecule of water consists of one hydrogen atom (with an electronegativity of 2.2) and one fluorine atom (with an electronegativity of 3.5). The more electronegative fluorine pulls the bonding electrons closer to itself resulting in the fluorine becoming slightly negative (δ-)and the hy-drogen becoming slightly positive(δ+). This means that a molecule of hydrogen fluorine contains a polar covalent bond.

A molecule of water consists of two hydrogen atoms (with an electronegativity of 2.2) and one oxygen atom (with an electronegativity of 3.5). The more electronegative oxygen pulls the bonding electrons from both of its bonds with the hydrogen atoms towards itself resulting in the oxygen becoming slightly negative (δ-) and the hydrogen's becoming slightly positive (δ+). This means that a molecule of water contains 2 polar covalent bonds.

Non-polar (pure) covalent bonds Pure (non-polar) covalent bonding occurs in molecules where there is little or no difference in the electronegativity of the bonding elements. In other words the electrons within the covalent bond are shared evenly. The diagram opposite shows the evenly shared electrons in a molecule of Fluorine. In the data booklet (page 11) fluorine is quoted as having an electronegativity of 4.0. Within a molecule of fluorine (F2) both atoms have the same electronegativity.

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Polar/non-polar molecules Deciding whether or not a covalent molecule is polar or pure requires us to look at 2 factors:

• The difference in electronegativities of the 2 bonding elements

• The shape/symmetry of the molecule Some molecule containing polar bonds like hydrogen fluoride and water end up with an overall polarity because the polar bonds are arranged asymmetrically within the molecule. Both molecules have a positive end (pole) and a negative end (pole). Molecules that have fixed poles are said to have permanent dipoles. Ammonia is an example of a polar molecule as it has a permanent dipole. However some molecules may contain many polar bonds but may not be polar molecules. In these instances the polar bonds are arranged symmetrically and thus cancel out the overall polarity. Tetrechloromethane is a common example of a molecule containing polar bonds, but exists as non-polar molecules.

The bonding continuum The bonding continuum seeks to place compounds on a scale in an attempt to make it easier to label the compound as either ionic, polar covalent or pure covalent. The bonding continuum places ionic compounds at one of the scale with pure covalent at the opposite end. Polar covalent compounds are placed in the middle. It is not perhaps surprising that the compounds with the greatest differences in electronegativity would be expected to be ionic and those will little or no difference in electronegativity to be pure covalent. It is along these principals that the bonding continuum has been created. Whilst the bonding continuum is good at differentiating between whether or not a compound contains pure covalent or ionic bonding it is less effective at distinguishing between polar covalent and ionic compounds as too often the differences in electronegativity are too close. At National 5 we learned that it is not simply enough to assume that because a compound contains both a metal and non-metal that the substance will be ionic as this is not always the case, e.g. tin(IV) iodide is covalent molecular.

Note - the classification of bonds into covalent (non-metals) and ionic (metals and non-metal) is an oversimplification. The bonding in most compounds is an intermediate between covalent and ionic.

The properties of a substance give a better indication of the type of bond present.

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Intermolecular forces and properties of compounds As the gaseous elements can be liquefied on cooling, e.g. diatomic hydrogen and oxygen and the monatomic noble gas-es, there must be forces of attraction between the molecules and the atoms of the noble gases. If such forces did not exist the above elements would be gases at all temperatures. The intermolecular forces of attraction are called London dispersion forces. London Dispersion Forces London dispersion exist between all non-polar molecules and atoms of the noble gases. This makes them the most significant intermolecular attraction that exists between non-polar molecules and atoms of the noble gases. As electrons are constantly moving within an atom or molecule, at any particular moment the electron distribution is unlikely to be evenly spread. Momentarily there may be more negative charge on one side of the molecule or at-om than the other. This results in one side of the particle having a momentarily slight positive charge and the other side a slight negative - this is known as a temporary dipole. London dispersion forces result from the attraction between the positive end of one dipole and the negative end of the other. The strength of the London dispersion depends on the size and mass of the molecule. The Greek letter delta (δ) is used to represent that the charge is only slight. The fact that London dispersion forces form between temporary dipoles means they are weak and easily broken. This means molecules containing London dispersion forces as there most significant intermolecular force tend to be liquids and gases at room temperature, e.g. butane is a gas at room temperature with a boiling point of 0⁰C. The temporary nature of the dipoles also means they are insoluble in polar solvents like water as they are unable to break the stronger permanent dipole-permanent dipole interaction holding the molecules of water together. Permanent dipole - Permanent dipole interactions Permanent dipole - permanent dipole interactions are the electrostatic attraction that exists between polar molecules containing permanent dipoles. Permanent dipole - permanent dipole interactions are stronger than London dispersion forces resulting in polar mol-ecules having higher melting and boiling points than non-polar molecules of similar size and mass (remember the strength of the London dispersion force depends on the size and mass of the molecule). An example that provides evidence of the increased strength of a permanent dipole - permanent dipole interactions compared to a London dispersion force would be when we compare propanone (C3H6O) and butane (C4H10). Both have the same molecular mass of 58 amu, but propanone boils at 56⁰C and butane at 0⁰C. Polar molecules are often soluble (miscible) in polar solvents because of the attraction between the polar intermolecular bond. Molecules containing the groups - O-H, N-H and the molecule hydrogen fluoride (HF) are so polar that the intermo-lecular force is given a special name - the Hydrogen bond.

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Hydrogen bonding Hydrogen bonding is a type of intermolecular force that exists between highly polar molecules, containing the N-H, O-H and the H-F bonds. They are considerably stronger than London dispersion forces and are formed between polar molecules containing a hydrogen atom which is bonded to an atom with a high attraction for electrons (electronegativity), i.e. fluorine, oxygen and nitrogen. Consider hydrogen fluoride, HF. For a non-metal atom hydrogen has a very low electronegativity (attraction for the bonding electrons). Fluorine on the other hand has a very high attraction for the electrons. This results in a highly polar covalent bond: Relatively string forces of attraction between the hydrogen atoms of the hydrogen fluoride molecules and the fluorine atoms of the neighbouring molecules will be set up. Hence, in the liquid and solid states hydrogen fluoride forms long bonded chains: The following substances will contain hydrogen bonds between their molecules N-H compounds: Ammonia (NH3) and amines, e.g. methyl amine (CH3NH2) O-H compounds: Water (H2O), alcohols, e.g. ethanol (C2H5OH) and carboxylic acids. H-F - Hydrogen fluoride. The presence of hydrogen bonding in compounds has a significant influence on their physical properties - boiling point, viscosity, surface tension and ice floating on water. Properties - Boiling points The graph opposite shows the plots of the boiling points of the hydrides of the elements in groups 4, 5, 6, 7 of the Periodic Table. The decrease in group 4 from stannane (SnH4) to methane (CH4) is expected as the strength of the London dispersion force decreases with decreasing molecular mass. The graph shows a similar and expected decrease from hydrogen telluride (H2Te) to hydrogen sulfide (H2S) but water (H2O) has a much higher boiling point than expected. A similar pattern occurs in the hydrides of groups 5 and 7 with ammonia (NH3) and hydrogen fluoride (HF) having a far higher boiling point than expected. In water liquid ammonia and hydrogen fluoride there must be present a relatively strong intermolecular attraction to account for the unexpected high boiling points. The relatively strong intermolecular attractions are called hydrogen bonds. Extra energy is required to break the hydrogen bonds between molecules and thus the boiling points are significantly higher.

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Properties - Viscosity The thickness of a liquid (its viscosity) is also increased by hydrogen bonding. Liquid covalent compounds containing London dispersion forces are non-viscous as the bonds are easily broken, e.g. benzene, petrol and tetrachloro-methane. Liquids containing hydrogen bonds between the molecules are more viscous as the molecules are held together more tightly, e.g. amines and the compounds containing the hydroxyl functional group (OH) like water, alcohol and carboxylic acids. Properties - Surface Tension In the above liquids hydrogen bonded molecules will be held together quite tightly on the surface of the liquid creating a sort of ‘skin effect’ on the surface, known as surface tension. The surface tension of water is high enough for small but relatively dense articles, such as razor blades and pins to float on the surface. Pond skaters are able to walk on water as the surface tension is able to support their mass. Properties - Density of ice Normally as liquids cool the density increases as the molecules move closer together as in the case of water to 4⁰C. But between 4C and 0C the density decreases with the formation of ice. This means below 4C the water molecules start to move apart - hydrogen bonding starts to order the molecules to give an open ad rigid structure. The fact that ice is less dense than water means that ponds and rivers will freeze from the surface downwards and the layers of ice insulates the water below, thus preventing complete solidification and ensuring the survival of fish, aquatic plants and other water-living creatures.