IGCSE Chemistry - Web viewIGCSE Chemistry. From the Edexcel ... This uses up the oxygen. On cooling, around 79cm3 of gas is left in the syringes ... Plastic bottles. Polypropene. Ropes.

IGCSE Chemistry

IGCSE Chemistry

From the Edexcel IGCSE 2009 Syllabus including triple science statements

CGPwnedWhen Chemistry and CGP books get pwned…

Unit 1: The Periodic TableThe periodic table contains about a hundred or so elements that have been currently discovered. The rows are known as periods and elements of the same period have the same number of electron shells. The columns are known as groups and elements of the same group have the same number of electrons on their outer shell. Group one has one outer electron,

An element is a substance that cannot be broken down into anything simpler. KCl for example (potassium chloride) is NOT an element because it can be broken down into K (potassium) and Cl (chlorine). The potassium and chlorine are the elements.

A compound is two or more elements chemically bonded together. An example would be KCl (potassium chloride), which consists of the elements potassium and chloride chemically bonded together.

Atoms are the building blocks of substances.

Molecules are two or more atoms bonded together. It doesn’t have to be a compound. Elements such as O2 and Br2 are diatomic molecules – they exist in pairs.

Atomic StructureAtoms are made up of protons, neutrons and electrons. Protons are positively charged. Electrons are negatively charged. Neutrons don’t have a charge.

An atom consists of a nucleus, which contains protons and neutrons; and some electron shells which surround the nucleus and contain electrons. The neutrons however, are different. The number of protons and the number of neutrons add up to make the mass number of an element.

Understanding the Lack of Reactivity in Noble Gases (Group 0)

Noble gases have eight electrons on their outer shell, therefore, there is no need for them to gain or lose electrons. Basically they have a full outer shell so they don’t need to react. This is what makes them so unreactive.

How to Read Each Square on The Periodic Table

You probably already know that the periodic table is made up of lots and lots of squares, each containing an element and information about it.

Anyways we already know what the atomic mass number is (the number of protons + neutrons). It says 12.011 here but this is probably because this picture came from some super complicated periodic table. In IGCSE level however, the atomic mass should read 12. Anyways, the atomic number is the number of protons (and electrons), so to find the number of neutrons, if asked to, simply subtract the atomic mass by the atomic number.

Example: Calculate the number of neutrons Carbon has.

The answer: 12 – 6 = 6 neutrons

The Arrangement of Electrons

Atoms are surrounded by electron shells which contain electrons. But the arrangement is the same for ALL the elements, not matter how different they are.

Each shell can only hold a certain number of electrons. The very first shell can hold only two electrons. The second shell can hold eight. The third sometimes appear full with eight but can expand to a total of eighteen. However, this is beyond GCSE level, and for now, the shells only hold eight.

So how do you find the electron configuration? Well let’s use potassium (K) as an example. Look up the atomic number of potassium. It should say 19. This tells you the number of protons, which is equal to the number of electrons so we can use that.

Arrange the electrons in shells, always filling up the inner shell before you go to the outer one. Remember the first, innermost shell can only take 2 electrons, the second one can take 8, and the third one, 8. You will find that you have one electron left. That goes on the fourth shell.

Your electron configuration should look like this: 2, 8, 8, 1.

Example: Work out the electron configuration of chlorine.

Chlorine has an atomic number of 17 – so 17 electrons. 17 – 2 (as the innermost shell only holds two electrons) = 1515 – 8 (as the second shell only holds eight electrons) = 7 (This number is the number of electrons Chlorine has on its outer shell).7 electrons does not fill up the third shell so we are left with the configuration: 2, 8, 7.

Isotopes

The number of neutrons in an atom can vary slightly. For example, there are three kinds of carbon atom, called carbon-12, carbon-13 and carbon-14. They all have the same number of protons, but the number of neutrons vary. These different atoms of carbon are called isotopes. Isotopes are atoms that have the same atomic number, but different mass numbers. They have the same number of protons, but different numbers of neutrons. The fact that they have varying numbers of neutrons makes no

difference whatsoever to their chemical reactions. The chemical properties are governed by the number and arrangement of the electrons.

Calculating Relative Atomic Mass (R.A.M.)Lets start this off with an example!

Example: Naturally occurring silver is 51.84% silver-107 and 48.16% silver-109. Calculate the relative atomic mass of silver.

r.a.m. (Ag) = (51.84/100 x 107) + (48.16/100 x 109)= 55.469 + 52.494=107.96

Now what did we do there? Well I simply calculated 51.84% of 107 (of silver) and 48.16% of 109 (of silver), and added the two answers! What we end up with is 107.96. Round that up to a whole number and the average relative atomic mass of silver is 108.

Calculating the Abundance (percentage) of an IsotopeExample: Copper consists of two isotopes, copper-63 and copper-65. Its relative atomic mass is 63.62. Find the abundance of each isotope.

Let y/100 = abundance of copper-63Let (100-y)/100 = abundance of copper-65

63.62 = (y/100 x 63) + [(100-y)/100 x 65]63.62 = 63y +6500 – 65y-2y = -135y = 69

Abundance of copper-63 = 69%Abundance of copper-65 = 100 – 69 = 31%

About Metals and Non-MetalsThe IGCSE spec. states you have to recall the positions of metals and non-metals on the periodic table. That’s easy! Its on page two. Have a look. Its colour-coded.

Anyways, this section covers 2.2, 2.3 and 2.5.

Metals

Metals tend to be shiny. They tend to have high melting and boiling points because of powerful attractions. Metals conduct heat and electricity because delocalized electrons are free to move throughout the structure. Metals are usually easy to shape due to their regular packed molecules. Metals react with water to form bases, and their oxides are also bases. They are good reducing agents because they lose electron.

Non-Metals

Non-metals tend to be brittle. They are poor conductors of heat and electricity. They form acidic oxides and are good oxydising agents because they gain electrons.

Aluminium Oxide

Aluminium oxide is amphoteric . It can neutralize both an acid and a base.

Reaction with acids

Aluminium oxide contains oxide ions and so reacts with acids in the same way as sodium or magnesium oxides. That means, for example, that aluminium oxide will react with hot dilute hydrochloric acid to give aluminium chloride solution.

In this (and similar reactions with other acids), aluminium oxide is showing the basic side of its amphoteric nature.

Reaction with bases

Aluminium oxide has also got an acidic side to its nature, and it shows this by reacting with bases such as sodium hydroxide solution. Various aluminates are formed - compounds where the aluminium is found in the negative ion. This is possible because aluminium has the ability to form covalent bonds with oxygen.

Group 1: The Alkali MetalsAlkali metals are metals that are part of group one. They are extremely reactive metals, and reactivity increases DOWNWARDS – in other words, lithium is the least reactive and francium.

Some Basic Physical Properties

Metal Melting Point (0C)

Boiling Point (0C) Density (g/cm3)

Lithium 181 1342 0.53Sodium 98 883 0.97Potassium 63 760 0.86Rubidium 39 686 1.53Francium 29 669 1.88

You can see that as reactivity increases, the melting and boiling points decreases; however, density increases. These points are very low for metals. Remember that potassium, sodium and lithium would float on water due to their densities. But why are they so reactive? Well they only have one electron to lose!

The metals are also very soft and easy to cut, becoming softer as you go down the group. They are shiny and silver when cut, but tarnish within seconds on exposure to air.

Storage and Handling

All these metals are extremely reactive. Anyways the metals will quickly react with air to form oxides, and react between rapidly and violently with water to form strongly alkaline solutions of metal hydroxides.

To stop them reacting with oxygen or water vapour in the air, lithium, sodium and potassium are stored under oil. Rubidium and caesium are so reactive that they have to be stored in sealed glass tubes to stop any possibility of oxygen getting at them.

Great care must be taken not to touch any of these metals with bare fingers. There could be enough sweat on your skin to give a reaction producing lots of heat and a very corrosive metal hydroxide.

Reactions with Water

All these metals react with water to produce a metal hydroxide and hydrogen.

Metal + Water Metal Hydroxide + Hydrogen

All the hydroxides are bases and turn pH paper purple.

With Sodium

The sodium floats because it is less dense than water. It melts because its melting point is low and a lot of heat is produced by the reaction. Observations would be that the sodium would turn into a ball and whiz around the surface of the water. It may form a white trail which is sodium hydroxide. This dissolves to make a strongly alkaline solution with the water. When lit, it produces a yellow flame.

With Lithium

The reaction is very similar to sodium’s reaction, except it is slower. The lithium does not melt due to its higher melting point. When lit, it produces a red flame.

With Potassium

Potassium’s reaction is faster than sodium’s. Enough heat is produced to ignite the hydrogen, which burns with a lilac flame. The reaction often ends with the potassium spitting around.

With Rubidium and Caesium – The Two Baddies

The reaction is so violent it can be explosive. When lit, Rubidium forms a red flame and Caesium forms a blue flame.

Explaining the Increase in Reactivity

The differences between reactions depend in part on how easily the outer electron of the metal is lost in each case. That depends on how strongly it is attracted to the nucleus. The more electron shells an atom has, the less powerful the attraction forces are. For example, Lithium is a lot less reactive than Potassium. This is because there are less shells which shield the full attraction of the nucleus from the This makes the electron harder to lose. However, potassium has a lot more electron shells which shield the outer electron from the nucleus. This weakens the attraction in compared to lithium, and therefore, the electron is easier to lose.

Compounds of Alkali Metals

All group one metal ions are colourless. That means that their compounds will be colourless or white unless they are combined with a coloured negative ion (remember metals would become positive ions because they lose electrons, whereas, most non-metals gain electrons). Potassium dichromate is orange, for example, because the dichromate ion is orange. Group one compounds are typical ionic solids and are mostly soluble in water.

Alkali Metals: Quick Notes

Group One so +1 charge One electron on outer shell

Reactivity increases downwards Density increases downwards Melting and Boiling points both decrease downwards Very soft and tarnish quickly in air Li, Na and K are stored under oil, whilst Rb and Cs are stored

in sealed glass tubes Reacts with air to form oxides Reacts with water to form alkaline hydroxides, which turns

pH paper purple Positive ions are formed and they are colourless Flame Colours: lithium, red; sodium, yellow; potassium, lilac;

rubidium, red; caesium, blue. Forget about Francium you don’t need to know much about

it.

Group 2: Alkali Earth MetalsAlkali earth metals belong to Group two. They are beryllium, magnesium, calcium, strontium, barium and radium. These metals are harder than those in group one. They are silvery grey in colour. They tarnish quickly, however they don’t just disappear into thin air because the oxides the metals form when reacting with air would form an outer coat that protects the metal from the air. They are good conductors of heat and electricity. They burn in oxygen to form white oxides. They react with water to form hydroxides and hydrogen, but the reaction is a lot less than that of group one. Also, reactivity increases down the group.

Flame Colours

-Calcium –brick red-Strontium -crimson-Barium -apple-green

Well that’s it for Group two!

Alkali Earth Metals: Quick Notes

Harder than group one metals Two electrons on outer shell (2+ charge) Form white oxides Forms hydroxides and hydrogen when reacting with water. Reaction is less vigorous than that of

group one Reaction increases downwards Silvery-Grey Flame Colours: calcium, brick red; strontium, crimson; barium, apple-green.

Group 7: The HalogensHalogens are group seven elements. Their elements are diatomic molecules. They exist in pairs, such as F2 and Cl2. These two elements are gases, bromine is a liquid and iodine is a solid. Astatine is radioactive.

These vapours and gases are poisonous. All these elements need to be handled in a fume cupboard.

Element State Colours

Flourine Gas Yellow

Chlorine Gas Green

Bromine Liquid Orange – Brown vapour

Iodine Solid Dark Grey – Purple vapour

Reactions with Hydrogen

The halogens react with hydrogen to form hydrogen halides such as hydrogen fluoride and hydrogen chloride. These are all steamy, acidic and poisonous gases. They are very soluble in water, reacting with it to produce solutions of acids. However as a gas, it is NOT an acid.

Reaction Between Sodium and ChlorideSodium burns in chlorine to produce the white solid sodium chloride – or salt!

2Na(s) + Cl2(g) 2NaCl(s)

In this reaction, sodium has been oxidized since it has lost electrons. Chlorine has been reduced.

Displacement Reactions with HalogensFinding the reactivity of halogens are done by reacting the elements with potassium halides. Colour change will indicate a reaction.

Note: Colour changes are due to the element being displaced. For example, the colour change from yellow to brown when chlorine reacted with potassium bromide was due to the fact that the bromine was displace. It was the brown of the Bromine that turned the solution brown.

Potassium is only a spectator ion. It does not change.

But now we have a problem. To distinguish whether bromine or iodine has been displaced is difficult, as both elements produce very similar shades of brown. What do we do? We add an organic solvent such as Volasil. When Volasil is added, the iodine turns pink while the bromine stays brown. Pretty neat huh?

These reactions are known as redox reactions, where oxidation and reduction are occurring (not just one of them).

Explaining the Trend in Reactivity of Halogens

As you go down the group, the oxidizing ability of the halogens falls due to the decreasing reactivity. When a halogen oxidizes something, it does so by removing electrons from it. Chlorine is a strong oxidizing agent because its atoms readily attract an extra electron to make chloride ions. Bromine is less successful. Why? This relates to electron shells again. In Chlorine, there are three shells which shield the nucleus’ attraction force from attracting another electron to gain a full outer electron shell. Bromine however, has a lot more shells to shield the attraction, therefore, the force is much weaker.

Halogens: Quick Notes

Diatomic molecules Seven electrons on outer shell Highly reactive – only need one electron to fill outer shell Form hydrogen halides when reacting with hydrogen

Halogen Reaction with HydrogenFlourine Violent explosion, even in the cold and darkChlorine Violent explosion if exposed to a flame or

sunlightBromine Mild explosion if a bromine

vapour/hydrogen mixture is exposed to a flame

Iodine Partial reaction to from hydrogen iodide if vapour is heated continuously with hydrogen

Observations

Potassium Chloride

Potassium Bromide

Potassium Iodide

Chlorine Yellow to Brown

Yellow to Brown

Bromine Brown Brown to Dark Brown

Iodine Brown Brown

Reaction increases as you go up the group Halogens can displace each other Volasil turns iodine pink

The Difference Between Hydrogen Chloride and Hydrochloric AcidHydrochloric acid is basically a solution of hydrogen chloride gas in water.

The Bronsted-Lowry Theory

Bronsted and Lowry defined acids and bases as the following:

-An acid donates a proton.-A base accepts a proton.

How is this related? Well, when hydrogen loses its only electron, it becomes a hydrogen ion (H+). In other words, it is also a proton, because it has lost all of its electrons (it only has one remember?).

When hydrogen chloride dissolves in water, a proton (the hydrogen ion) is transferred to the water. This gives us the equation:

H2O(l) + HCl(g) H3O+(aq) + Cl-(aq)

The H3O+ ion is called a hydroxonium ion. We normally write it as H+(aq). You can think of it as a hydrogen ion riding on a water molecule.

So in this example, HCl is an acid because it donates a proton (the hydrogen ion) to water.

So the real differences? Hydrogen chloride is NOT an acid and is a gas. Hydrochloric acid is an aqueous solution of hydrogen chloride.

Hydrogen Chloride and Methylbenzene

Explaining Water Being a Polar Molecule

Water is a polar molecule. Electrons in water are attracted towards the oxygen end of the bond, which leaves it slightly negative. This leaves hydrogen slightly short of electrons, and

therefore, making it slightly positive, just like the picture to the left. Because of this electrical distortion, water is described as a polar molecule.

When something such as sodium chloride is being dissolved in water, the slightly positive hydrogens cluster around the chlorine, whereas, the slightly negative oxygen cluster around the sodium. The water molecules then literally pull the sodium chloride crystal apart.

This pull doesn’t work on every molecule. Magnesium oxide isn’t soluble in water because the water molecules aren’t strong enough to break the magnesium-oxygen attractions.What’s so special about methylbenzene?

Well, methylbenzene is not a polar molecule. It is unable to pull the hydrogen and chlorine apart and therefore, hydrochloric acid won’t be formed.

Oxygen and Oxides

(2.15) Composition of Air

This is the approximate composition of air. Memorize it.

There are also very small amounts of noble gases in the air.

(2.16)Showing That Air Contains About 1/5 Oxygen

Using Copper

The apparatus originally contains 100cm3 of air. This is pushed backwards and forwards of the heated copper, which turns black as copper(II) oxide is formed. This uses up the oxygen. On cooling, around 79cm3 of gas is left in the syringes – 21% has been used up. Therefore, the air contains 21% of oxygen.

Using the Rusting of Iron

Iron rusts in damp air, using oxygen up as it does so. The experiment shows some damp iron wool in a test tube containing air. The tube is inverted in a beaker of water and the level of the water in the tube is marked by a rubber band. The tube is left for a week or so for the iron to use up the oxygen to make…guess…iron oxide!

The water level rises in the tube as the oxygen is used up, and the new level can be marked using a second rubber band. You can find the actual volumes of the gases at the end of the experiment by filling the tube with water to each of the rubber bands in turn, and pouring it into a measuring cylinder. If the original volume was, say, 15cm3, and the final volume was 12cm3, then the oxygen used up measures 3cm3.

The percentage of oxygen in air was 3/15 x 100 = 20%.

Burning Phosphorus

This can be done by putting a bell jar into a beaker filled with water. Phosphorus on an evaporating dish is placed onto the water (the jar has no bottom). It is then touched with a hot metal rod, which starts the reaction between phosphorus and oxygen. Phosphorus uses up the oxygen to form phosphorus oxide, lowering pressure in the jar and therefore, making water levels rise in the jar. The water should rise up by 20%.

Equation:

(2.17) Making Oxygen in the Lab

Oxygen is most easily made in the lab from hydrogen peroxide solution using manganese(IV) oxide as a catalyst. The reaction is known as the catalytic decomposition (splitting up using a catalyst) of hydrogen peroxide.

Gas Amount in Air (%)Nitrogen 78.1Oxygen 21.0Argon 0.9Carbon Dioxide 0.04

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

Reaction of Oxygen with Magnesium, Carbon and Sulfur

Magnesium Magnesium reacts with oxygen to produced white, powdery magnesium oxide. It produces a bright white flame during the reaction. It is a base.

2Mg(s) + O2(g) 2MgO(s)

With Sulfur Sulfur burns in oxygen with a tiny blue flame. Poisonous, colourless sulfur dioxide is produced. It is an acidic oxide.

S(s) + O2(g) SO2(g)

With Carbon Carbon burns in oxygen if heated strongly to give colourless carbon dioxide. Depending on the purity of the carbon, a small yellow-orange flame may be produced.

C(s) + O2(g) CO2(g)

Carbon DioxidePreparing It in the Lab

Carbon dioxide is made by the reaction between dilute hydrochloric acid and calcium carbonate in the form of marble chips.

CaCO3(s) + 2HCl(aq) CaCl2(aq) + CO2(g) + H2O(l)

Formation of Carbon Dioxide from Thermal Decomposition of Metal Carbonates

Key thing here: When heating metal carbonate, you get:

Metal Carbonate Metal Oxide + Carbon Dioxide

Here is the picture of the experiment setup:

Properties of Carbon Dioxide

Colourless gas, denser than air, slightly soluble in water

Used in carbonated (fizzy) drinks because it dissolves in water under pressure. When bottle is opened, pressure falls and gas bubbles out of the solution.

Used in fire extinguishers to put out electrical fires, or those caused by burning liquids, where using water could cause problems. The carbon dioxide sinks onto the flames and prevents any more oxygen from reaching them.

Turns limewater cloudy white (limewater is calcium hydroxide – work out the equation yourself – water is one of the products).

Carbon Dioxide and Sulfur Dioxide – Their Reactions With WaterCarbon Dioxide Carbonic acid is produced when

carbon dioxide reacts with water. It is a weak acid. This reaction can be reversed by simply heating or boiling the acid.

CO2(aq) + H2O(l) ⇌ H+(aq) + HCO3–

(aq)

Sulphur Dioxide Sulfur dioxide reacts with water to H2O(l) + SO2(g) H2SO3(aq)

form a weak acid known as sulfurous acid,

Sulfur Dioxide, Nitrogen Oxide and the EnvironmentAcid rain is caused when oxygen and water in the atmosphere react with sulfur dioxide to produce sulfuric acid (ouch), or with various oxides of nitrogen to give nitric acid. These mainly come from power stations, burning fossil fuels, motor vehicles etc.

Acid rain can kill trees and make lakes so acidic it cannot support life. Limestone and some metals such as iron are also attacked by acid rain.

The solution to acid rain involves removing sulfur from fuels, using catalytic converters in cars and scrubbing the gases from power stations to remove the oxides. The catalyst helps convert nitrogen oxides into harmless nitrogen gas but has no effect on sulfur dioxide.

Methods of Separation

Filtration: For separating an insoluble solid from a liquid, or a soluble solid from an insoluble one.

Sand can be separated from water by pouring the mixture down a funnel with filter paper. The sand will collect at the filter paper.

It can also be used to separate sand from something like salt by dissolving salt in water (which leaves you with sand mixed with salt water). The mixture can then be poured down a funnel. The sand that collects at the top should can be rinsed and dried. The water can be evaporated from the salt by heating with a Bunsen burner. Back

Chromatography: For separating liquids by dissolving them in a solvent. The dyes that make up the ink differ in two important ways:

How strongly they stick to the paper How soluble they are with the solvent

An example would be separating ink colours or plant dyes. A dot of the ink/dye would be drawn onto a piece of paper. It would then be left in water, which acts as a solvent. Because different colours have different solubility levels, some colours would travel up further on the paper.

Crystallization: Mainly used for purifying substances by forming crystals from a precipitating solution. Crystallization refers to the forming of solid crystals from a homogenous (solution) mixture.

An example would be forming pure salt crystals. This is done by dissolving the

impure salt into a solvent such as water. The salt solution is then allowed to cool. As it does, pure salt crystals would form at the bottom of the water, whereas, the impure substances would be left in the water. The crystals can then be rinsed with a chilled solution and dried.

Distillation: Distillation is good from separating a liquid from a solution.

An example would be separating water from a salt solution. The solution would be heated at the liquid’s boiling point, in this case 1000C, so it will leave the solution as a vapour. The vapour would then condense into a liquid with the help of the cooling water. The vapour, now as a liquid, would fall into the beaker.

Fractional Distillation: Fractional Distillation is used to separate two liquids based on their boiling points.

An example would be separating ethanol from water. Ethanol has a lower boiling water than water (at about 780C), therefore, the heating is monitored (using the thermometer) to ensure that the temperature does not reach 1000C (the boiling point of water). Anyways, the ethanol would turn into a vapour and travel out of the flask. It would then condense into its liquid form with the help of the cooling water and fall into the beaker.

Unit 2: Structure and Bonding

Ionic Bonding

Ionic bonding is the bonding in which there has been a transfer of electrons from one atom to another to produce ions. The substance is held together by strong electrostatic attractions between positive and negative ions. Ions are formed when it gains or loses electrons. Ones that gain forms negative ions, and ones that lose form positive ions.

A positive ion is called a cation. A negative ion is called an anion.

You can find the charge of an ion by looking at the group it belongs to. If it belongs to groups 1-4, it has a charge of 1-4+ (they are positive), whereas, if it belongs to groups 5-0, it has a charge of 3—0. Below is a table containing charges of common ionic compounds and transition metals.

This is an example of a dot and cross

diagram. The crosses represent the electrons on the

Ion Symbol Charge

Silver Ag 1+

Copper (I) Cu 1+

Ammonium NH4 1+

Copper (II) Cu 2+

Cobalt Co 2+

Nickel Ni 2+

Zinc Zn 2+

Iron (II) Fe 2+

Chromium Cr 3+

Iron (III) Fe 3+

Ion Symbol Charge

Hydroxide OH 1-

Nitrate NO3 1-

Hydrogen Carbonate

HCO3 1-

Carbonate CO3 2-

Sulphate SO4 2-

Phosphide P 3-

Phosphate PO4 3-

Nitride N 3-

sodium (anion) and the dots represent the electrons on the chlorine (cation). In a dot and cross diagram, you must use arrows to show which electrons are moved from the anion to the cation. On the final diagram, you mark the new electron(s) on the cation as a cross.

Boiling and Melting Points of Ionic Compounds

Ionic compounds have high boiling and melting points due to strong intermolecular forces between the atoms. This is because when the ions are formed during an ionic reaction, one of them would be positive, and one would be negative. Positive and negative attract and therefore, you get something like a strong magnet.

As ionic charge increases, so does the melting/boiling points. Ions with 2+ and 2- would have stronger attraction because their charges a stronger, whereas, ions with 1+ and 1- would still have a strong attraction, but less stronger than 2+- compounds.

Structure of Ionic Compounds

An ionic crystal consists of giant three-dimensional lattices held together by strong electrostatic attractions between the positive and negative ions.

Structure of Sodium Chloride

This is the basic structure of a sodium chloride crystal. The green is the chloride and the blue is the sodium. Remember that each sodium is touched by six chlorides and each chloride is touched by six sodiums. Look at the middle atoms if unclear. Remember, this structure repeats itself over and over.

Ionic bonds always produce giant structures.Ions form closely packed regular lattice arrangement.They have high melting/boiling points.The crystals tend to be brittle.Compounds tend to be soluble in water and insoluble in organic solvents.

Covalent Bonding

Covalent bonding is formed by sharing a pair of electrons between two atoms. This is so that both atoms can achieve a full outer shell. It is a strong attraction between the bonding pair of electrons and the nuclei of the atoms involved. Covalent compounds are only formed when the reactants are non-metals.

Diagrams YOU Need to Know…

Element

Diagram Element

Diagram Element

Diagram

H2 CH4 CO2

Cl2 NH3 Ethane

HCl(g) O2

H2O N2 Ethene

Simple Molecular Structures

These are gases, liquids or solids with low melting points. Examples include water, chlorine, oxygen…etcThe covalent bonds between the atoms in a molecule are strong.However, the forces of attraction between these molecules (inter-molecular forces) are weak.They have low melting points, since not a lot of heat is needed to provide the energy for the molecules to move away from each other, hence, overcome the intermolecular forces between them.They tend to be insoluble in water.They are often soluble in organic solvents.They do not conduct electricity because the molecules have no overall charge and there are no electrons mobile enough to move from molecule to molecule.

Giant Covalent Structures

There are no charged ions. ALL the atoms are joined up to their adjacent atom by extremely strong covalent bonds and

packed into giant regular lattices.They have very high melting points, since a lot of heat is needed to provide the energy to break apart the many strong covalent bonds.They tend to be insoluble in water.They do not conduct electricity.

Diamond

The diamond is the hardest natural substance. It is a form of pure carbon. Each carbon atom forms four covalent bonds to the other carbon atoms. They are arranged in a tetrahedral arrangement. Diamond has a very high melting point, obviously due to very strong carbon-carbon bonds. It does not conduct electricity because all the electrons in the outer levels of the carbon atoms are tightly bonded between the atoms. None of them are free to move around. Diamond is insoluble – like, to both water and other solvents.

Use of Diamond

Saw blades can be tipped with diamonds in high-speed cutting tools used on stone and concrete. The strong tetrahedral structure makes the diamond hard, making it suitable for this purpose.

Graphite

Graphite is arranged differently – it has a layer structure. Each graphite layer is strong, but it is easy to separate individual graphite layers. Each carbon atom only forms three covalent bonds. Graphite conducts electricity because the fourth electron is free to move around.

Use of Graphite

Because of the layered structure, graphite can be used as a dry lubricant to lubricate locks.

Metallic Crystals

Metals are giant structures which consist of a regular array of positive ions in a sea of delocalized electrons. When metal atoms bond together to form solid, visible metal, their outer electrons are no longer attached to particular electrons and are free to move around the whole structure.

Metals are able to conduct electricity because the delocalized electrons are free to move throughout the structure. The energy is picked up by the electrons and moved around the metals, transferring the electricity throughout the whole structure. The same goes to heat energy.

Metals are easy to shape because their regular packing makes it simple for atoms to slide over each other. Metals are said to be malleable.

Introduction to Electrolysis

In metals and carbon, electricity and electric current is simply a flow of electrons or ions. Electrolysis is the chemical change caused by passing an electric current through a compound which is either molten or in a solution. An electrolyte is a substance that undergoes electrolysis. It contains ions. It is the movement of the ions, which are responsible for both the conduction of electricity and the chemical changes that take place. Covalent compounds are not electrolytes and don’t conduct electricity because they have no free moving electrons. Ionic compounds only conduct electricity when molten or in a solution because the ions separate and are free to move. These particles can then carry the electric current.

Diffusion

Diffusion happens when particles spread from higher to lower concentration. It requires a concentration gradient).

Experiment to distinguish between electrolytes and non-electrolytes

i. Dissolve substance in water, or if possible, melt it. ii. Put a conductivity tester into the substance.iii. If the light bulb lights up, it is an electrolyte.

Explain: When dissolved in water, free moving electrons are able to carry the electric current across from the cathode to the anode, completing the circuit and lighting the bulb. If the light bulb does not light up, the substance is obviously not an electrolyte.

But sugar dissolves, why does the bulb not light up? Sugar is a covalent structure.

Potassium Manganate (VII) Experiment

Diffusion through liquids is very slow if the liquid is totally still. This can be shown but dropping a piece of potassium manganate (VII) into water. It can take days for the colour to spread because the gap between each particle is small.The Bromine Experiment

Showing diffusion in gases can be done by filling a lower gas jar with bromine gas and topping it with a gas jar filled with air. The bromine particles and air particles will eventually bounce around to give an even mixture.

The Ammonium Chloride Experiment

This experiment is used to show that particles in different gases travel at different speeds. It relies on the reaction between ammonia and hydrogen chloride gases to give white solid ammonium chloride.

A white ring of ammonium chloride would form near the hydrochloric acid. This shows that ammonia particles have travelled further to reach the hydrogen chloride gas, showing that it travels faster.

Dilution

Dilution is the reduction of concentration in a solution.

Showing Dilution and Leading to the Idea of Small Sized Particles

Suppose you dissolve 0.1g of potassium manganate (VII) in 10cm3 of water to give a deep purple solution. Assume the smallest drop you can see is 1/1000cm3. The whole solution will be made up of 10000 drops, each drop containing 0.00001g of potassium manganate (VII).

Suppose you dilute this down 10 times by taking 1cm3 of the solution and making it up to 10cm3 with more water. Continue doing this until the colour is too faint to see. By the time of the fifth dilution, each drop will only contain a billionth of a gram of potassium manganate (VII). If you only needed one ‘particle’ of potassium manganate (VII) per drop in order to see the colour, the ‘particle’ can’t weigh more than a billionth of a gram.

IS this a good answer? Nowhere near it! A potassium manganate (VII) particle actually weighs about 0.00000000000000000000026g! In reality, you need huge numbers of particles in each drop in order to see the colour.

Don’t worry I don’t get this either…

Unit 3: Organic ChemistryOrganic chemistry is mainly based around hydrocarbons – compounds made only up of hydrogen and carbon. It is drawn with lines joining carbons and hydrogen. All carbon bonds have to be bonded to hydrogen – if not something else. The left picture below shows carbons with all bonds taken up (ethane). The right picture below shows an incorrect picture of a hydrocarbon because one of the carbons has a free bond.

Hydrocarbon – compounds that contain carbon and hydrogen only.

Homologous series – family of compounds with similar properties because they have similar bonding. They show a graduation in physical properties (mpt/bpt) and similar chemical properties such as the general formula. Alkanes are the simplest.

Saturated – when carbon cannot take anymore bonds – single carbon-carbon bonds.

Unsaturated – presence of a carbon-carbon double bond. General formula – The formula of different homologous series of carbons. Isomers – molecules with the same molecular formula but different structural formulae.

Learning the Code

Do you have to remember the formula for propane, butane, ethane…? No! You can work it out yourself! The first part of the name tells you how many carbons there are in the longest chain (not necessarily in total). By the way you have to learn these – at least the first five. It helps.

For example: propane (left) has three carbons. Butane (right) has four carbons.

Alkanes and Alkenes

Alkanes and Alkenes are two homologous series.

Pentene has a five carbon chain with a double bond.

IsomersWe know what isomers are.

Code Number of Carbons

Meth 1Eth 2Prop 3But 4Pent 5Hex 6Hept 7Oct 8Non 9Dec 10

Ending Meaning?‘ane’ All carbon bonds

are filled with hydrogen – i.e. they are saturated hydrocarbons.

‘ene’ There is a double carbon-carbon bond – they are unsaturated hydrocarbons.

Coding for Double Bonds

For things like pentene and butene, there are many places you can put the double bonds in.

Pent-1-ene means pentene with the double bond on the first carbon-carbon bond (right).Pent-2-ene means pentene with the double bond on the second carbon-carbon bond and so on…

But wait! What about pent-4-ene and pent-5-ene? Those don’t exist. Why? Because pent-4-ene is pent-2-ene flipped over, and pent-5-ene is pent-1-ene flipped over!

Flip: C-C-C=C-C and you get C-C=C-C-C!

Methyl and Ethyl Groups

If the hydrocarbon has a methyl or ethyl group, these two come first, before the coding for the number of carbons in the chain. But before even the ‘methyl’ or the ‘ethyl’ there is a number and hyphen to show which carbon has the methyl or ethyl branch.

For example, this is 2-methylbutane

As you can see, there is a methyl group branching off the second carbon. The rules are similar to double bonds though, there is no such thing as 3-methylbutane because that is basically 2-methylbutane flipped over.

But wait! There are five carbons! Why isn’t it 2-methylpentane? Because remember, these names are based on the longest carbon chain in the hydrocarbon and the longest carbon chain there is 4, hence, butane. This means that 2-methylbutane is an isomer of pentane C5H12.

Some Isomers of Butane C4H10

Some Isomers of Pentane C5H12

Code MeaningMethyl Has a branch of

CH3 coming off one of the bonds.

Ethyl Has a branch of CH3CH2 coming off.

AlkanesAlkanes are a homologous series of saturated hydrocarbons. The first five are methane, ethane, propane, butane and pentane.

The general formula for alkanes is: CnH2n + 2

For example, ethane:

Ethane has two carbons, so n=2.The formula of ethane must be C2H2(2) + 2 = C2H6.

Complete Combustion of Alkanes

If there is enough oxygen, alkanes will burn in oxygen completely to give carbon dioxide and water. The general equation for combustion:

Hydrocarbon + Oxygen Carbon dioxide + Water

The combustion of methane would be: CH4(g) + 202(g) C02(g) + 2H20(l)

Note: Balancing combustion equations can be annoying. An easy way would be to balance them in the order of carbon, hydrogen then oxygen.

Incomplete Combustion

If there isn’t enough oxygen, you get incomplete combustion, in which carbon monoxide and water are produced instead. Carbon monoxide is a colourless, odourless and poisonous gas. It is dangerous because it can combine to our haemoglobin and stop it from carrying oxygen. As a result, you get ill or even die because oxygen cannot travel to all parts of your body.

Reaction with Bromine

Alkanes react with bromine under the presence of ultra-violet light. One hydrogen from the hydrocarbon would be replaced by a bromine atom. This is known as a substitution reaction. Bromine can be used as an indicator for alkanes and alkenes without UV light. Adding bromine water to alkanes produces no colour change. Reacting bromine water to alkenes make it turn from brown to colourless.

However, if the mixture of bromine and methane is reacted under UV light, it loses its colour, a mixture of bromomethane and hydrogen bromide gases is formed.

CH4(g) + Br2(g) CH3Br (g) + HBr(g)

AlkenesAlkenes have double bonds, making them unsaturated hydrocarbons.

Alkenes have the general formula of CnH2n – the first four being ethene, propene, butene and pentene.

Combustion

Like alkanes, alkenes burn in oxygen or air to give carbon dioxide and water.

Reaction with Bromine

Alkenes undergo addition reactions, in which part of the double bond breaks and is used to join other atoms onto the two carbon atoms. When added to alkenes, and the test tube is shook, the brown of the bromine would be decolourised, making it suitable as a test for alkenes.

The product of reacting ethene to bromine gives 1, 2-dibromoethane and is a colourless liquid.

CH2=CH2(g) + Br2(aq) CH2BrCH2Br(l)

EthanolAll alcohols contain an –OH group attached to a carbon chain. Ethanol is C2H5OH.

Production of Ethanol

Hydration of Ethene

Ethanol can be made by reacting ethene with steam (because it contains more energy) – a process known as hydration.

CH2=CH2(g) + H2O(g) CH3CH2OH(g)

Only a small portion of ethene reacts. The ethanol is condensed as a liquid and the unreacted ethene is recycled.Explaining the Choice of Temperature

Reversible reactions happen in two ways – while ethene is being converted into ethanol, ethanol is also being converted back into ethene. Reversible reactions can also shift the equilibrium – or ‘alter’ the reaction. Since the reaction is exothermic – the reaction produces lots of heat. If you increase the temperature, the reaction won’t like it because it is already producing heat, therefore, it would ‘adapt’ to the conditions by making more ethene so less heat will be produced. On the contrary, if you decrease the temperature, the reaction would ‘adapt’ to this by increasing back the temperature; by producing more ethanol – in other words, push the equilibrium to the favourable/forward reaction. However, making the temperature too low would mean super slow reaction, although more ethanol would be produced. 300 degrees is therefore, a compromise temperature producing an acceptable yield of ethanol in a short time.

Explaining High Pressure

In the equation, you have two moles (one mole of ethene and one mole of water) on the left, and one mole (of ethanol) on the right. Increasing the pressure would mean the equilibrium would be shifted forwards. Why? The reaction would ‘adapt’ to the conditions by producing more ethanol because you only get one mole of ethanol – which takes less space than two moles of ethene and water.

What is Needed:

Ethene and steam 3000C 60-70

atmospheres Phosphoric acid

catalyst

Also, there’s the collision theory. Increasing the pressure means that there’d be less space for the atoms to move. The atoms would also move with more force. This increases the frequency of collisions.

The problem: it’s expensive and ethene might polymerise and turn into polyethene.

Fermentation

Yeast is added to a sugar or starch solution at 30 0 C for several days in the absence of air for anaerobic respiration. Enzymes in the yeast lower the activation energy, increasing the rate of conversion of the sugar into ethanol and carbon dioxide. However, they first have to break the sugars into smaller sugars like glucose. In fact, ethanoic acid is produced and then converted into ethanol.

For example, sucrose:

C12H22O11(aq) + H20 C6H12O6(aq) + C6H12O6(aq) sucrose + water glucose + fructose

C6H12O6(aq) 2C2H5OH (aq) + 2CO2(g) glucose/fructose ethanol + carbon dioxide

The yeast then gets killed in the mixture, which means that the ethanol produced is impure. To purify it, the alcohol must undergo fractional distillation.

Comparing the two methods…

Fermentation HydrationUse of Resources Uses renewable resources – sugar

beet or sugar cane, corn and other starchy materials.

Uses non-renewable resources – once oil gets used up, they’re screwed.

Type of Process A batch process – everything is mixed and left for several days. It is then removed and a new reaction is set up – quite inefficient.

A continuous flow process – a stream of reactants is constantly passed over the catalyst – more efficient.

Rate of Reaction Slow, takes several days. RapidQuality of Product Produces impure ethanol that needs

further processing.Produces much purer ethanol.

Reaction Conditions Uses gentle temperatures and ordinary pressure – relying on anaerobic respiration of yeast.

Uses high temperatures and pressures, needing a high input of energy – expensive.

Common Question: Which method would poorer places like Brazil use and why? [3 marks]Answer: Fermentation, because Brazil has the weather conditions to grow large yields of sugar cane and they don’t have access to crude oil.

Dehydration of Ethanol into Ethene

Dehydration of ethanol produces ethene and water, using hot aluminium oxide as a catalyst.

CH3CH2OH(g) CH2=CH2(g) + H2O(l)

Crude OilCrude oil is a mixture of hydrocarbons. These chains can be super long or super short.

The Trend in Boiling Point and Viscosity

Viscosity means how runny something is Volatile means how easy it turns into vapour at room temperature

As the number of carbon atoms in molecules increases and gets bigger, intermolecular attractions also increase, making it more difficult to pull one molecule away from neighbouring ones. As they get bigger, these changes occur:

Boiling point increases – the larger the molecule, the higher the boiling point due to stronger intermolecular attractions.

Liquids become less volatile – the bigger the hydrocarbon, the more slowly it evaporates in room temperature. This is again, due to strong intermolecular attractions.

Liquids become more viscous (flow less easily) – Small hydrocarbons are runny, but large ones are much stickier and gooey (and viscous) because of intermolecular attractions.

Bigger hydrocarbons do not burn as easily, meaning they are less useful.

The Fractionating Column

Crude oil is separated in fractionating column. This process is fractional distillation, and splits crude oil into various fractions depending on their boiling points and size.

Note: forget about Naphtha

Fraction UsesRefinery gases A mixture of methane, ethane, propane and butane.

Commonly used for domestic heating and cooking.Gasoline CarsKerosene Used as fuel for jet aircraft.

As domestic heating oil.As ‘paraffin’ for small heaters and lamps.

Diesel oil For buses, lorries, some cars and railway engines.

Some is cracked to produce more petrol.Fuel oil For ships

Industrial heatingBitumen Residue from the bottom which can be used for roads.

Combustion and Incomplete Combustion

Combustion of hydrocarbons produces carbon dioxide and water – exothermic. Incomplete combustion of hydrocarbons produces carbon monoxide and water – in which

carbon monoxide is dangerous because it can bind to haemoglobin and prevent it from carrying oxygen.

In car engines, the temperature reached is high enough to allow nitrogen and oxygen from the air to react, forming nitrogen oxides. This contributes to smog and causes irritation to human mucus membranes. As well as that, nitrogen oxides can react with water in the atmosphere and from nitric acid – or acid rain.

Cracking

The Crude Oil Problem

Amounts of each fraction you get depend on the proportions of various hydrocarbons in the original crude oil. Far more petrol is needed, than something like bitumen. In other words, fractional distillation of crude oil produces more long-chain hydrocarbons than can be used directly, and fewer short-chain hydrocarbons than required.

The solution? Cracking! Cracking is a useful process in which large hydrocarbon molecules are broken into smaller ones. Most of the hydrocarbons found in crude oil are long-chain alkanes. Cracking can convert these into alkenes and shorter alkanes. It is an example of thermal decomposition.

How it Works

The fraction is heated to give a gas and is passed over a catalyst of silica or alumina with a temperature of 600-700 o C .

Long alkane alkene + alkane

Sometimes you may get more than one type of alkene/alkane. Make sure the numbers of carbon and hydrogen are balanced.

In an equation, this would read: hexane butane + ethene C6H14 C4H10 + C2H4

What is Needed:

Alumina/Silica as catalyst

600-7000C

PolymersAlkenes can be used to make polymers. Polymers are big long molecules of single units called monomers. Molecules containing carbon-carbon double bonds can be joined together. Part of the double bond is broken and used to join to other monomers. Joining up lots of monomers to make a polymer is called addition polymerisation.

How to Draw a Polymer

It’s like drawing a hydrocarbon, except the ends are left blank (so it can join more).

As for the repeating unit (which is the unit that keeps repeating itself), you show the alkene (or the monomer) with its double bond opened up. You then enclose it with brackets and put an ‘n’ to its right.

Polymers to Know

Polymer Repeating Unit How it looks together UsesPolyethene Plastic bags

Plastic bottles

Polypropene RopesCrates

Polychloroethene

PVC for drainpipes or windowsElectrical insulation

Nylon – Condensation PolymerIn condensation polymerisation, when two monomers combine, a small molecule such as water or hydrogen chloride is lost. Nylon is made through condensation polymerisation.

The two monomers that make up nylon:

Hexanedioic acid

From a family of compounds called dicarboxylic acids.

1,6-Diaminohexane

From a family known as diamines.

Joining them together…The lost of a Water Molecule

As a block diagram (where the (CH2)6 and (CH2)4 become blocks to make it look easier)

End Note: Sometimes you may be given ClOCCH2 CH2 CH2 CH2COCl instead of hexanedioic acid. In this case, just do the same thing, with the lost of hydrogen chloride HCl.

Unit 4: Analytical Chemistry and Kinestics

Tests for Ions and Gases

Flame Tests: Taking a piece of nichrome, make loop at the end and dip into salts containing ions

Ion ColourLi+ Crimson redNa+ Yellow orangeK+ LilacCa2+ Brick red/orange red

Using Sodium Hydroxide solution

Ion Colour of PrecipitateCu2+ BlueFe2+ Sludgy Green (or just green)Fe3+ Orange Brown (rust)

For Ammonium Ions (NH4+)

Heat gently and add sodium hydroxide solution. It will give off a distinctive smell of ammoni a (NH3). Ammonia can be tested by holding a damp red litmus paper . Since it is alkaline, it will turn damp red litmus paper from red to blue.

Using Dilute Nitric Acid and Silver Nitrate Solution

Ion Colour of PrecipitateCl- WhiteBr- Pale CreamI- Yellow

For Sulphate Ions (SO42-)

Using dilute hydrochloric acid solution and then adding barium chloride solution to form a white precipitate of barium sulphate.

For Carbonate Ions (CO32-)

Using dilute hydrochloric acid to react with the carbonate, to produce carbon dioxide gas which can be tested by bubbling through limewater, turning it from colourless to cloudy, milky white.

Tests for Gases

Gas Test ResultHydrogen Hold a lit splint in presence of hydrogen

gas.Produces a squeaky pop.

Oxygen Hold a glowing splint in presence of oxygen gas.

Glowing splint relights.

Carbon Dioxide

Bubble through limewater. Turns limewater from colourless to cloudy, milky white.

Ammonia Hold damp red litmus paper in ammonia gas.

Turns damp red litmus paper blue.

Chlorine Hold damp blue litmus paper in chlorine gas.

Bleaches or turns blue litmus paper white.

Solubility Patterns

All nitrates are soluble. All sodium, potassium and ammonium compounds are soluble. Most carbonates and hydroxides are insoluble except for sodium, potassium and ammonium. All sulphates are soluble except barium and lead(II) sulphate. All chlorides are soluble except lead(II) and silver chloride.

Reactions of Metals to Acids

Metals react very similarly to dilute hydrochloric acid and dilute sulphuric acid.

Metals Reaction to AcidMagnesium Rapid fizzing, mixture gets very hot, colourless magnesium sulphate/chloride

solution forms.Aluminium Is slow due to its coat of aluminium oxide which prevents aluminium from

contacting the acid. On heating, this layer is removed, aluminium will start fizzing rapidly – abit like Mg.

Zinc Zinc reacts slowly with cold dilute acid and may produce some effervescence. On heating however, it fizzes more.

Iron Iron also reacts slowly with cold dilute acid and will produce abit of effervescence when heated.

Combustion of Hydrogen

Bonds are broken in the hydrogen and oxygen molecules. These form new bonds of water molecules. This reaction is exothermic, and gives out water in the form of steam, before it condenses into a liquid. The reaction is:

2H2(g) + O2(g) 2H2O(l)

Testing for Water

Water turns white anhydrous copper(II) sulphate blue. It’s reaction is

CuSO4(s) + 5H2O(l) CuSO4•5H2O

Or you can use cobalt chloride paper – which turns pink in the presence of water.

You can check the purity of water by showing that it freezes at exactly 0ºC and boils at exactly 100ºC.

Rates of Reactions

Experiment Setup

To measure the effects of changes in surface area, concentration of solutions, temperature and use of catalyst, you can react calcium carbonate marble chips with dilute hydrochloric acid and measure the mass of CO2 produced by weighing the difference in mass of the reactants and the mass of the products (there won’t be any change in mass produced, because the initial mass of reaction will equal the final mass, however, since carbon dioxide gas is formed, this will escape from the flask, and therefore, the amount of mass lost will be the mass of carbon dioxide produced. Plot the results on a graph with mass against time and you’ll get an upward curve.

For the reaction to occur, acid particles must collide with the surface of the marble chips. As the acid particles get used up, the collision rate decreases, so the reaction slows down.

Changes in Surface Area of Solid

You can repeat the above experiment by keeping the same mass of marble chips, just using smaller ones to increase the surface area. The reaction happens faster. You have to remember the graphs. Notice however, that in the end, the amount of carbon dioxide produced is still the same – just that the small chips experiment happens faster.

Why does it happen faster? Because the surface area in contact with the gas or liquid is much greater. Less marble chip particles are hidden away from the acid particles.

Changes in the Concentration of Solutions

Repeat the original experiment but using hydrochloric acid only half as concentrated as before. The graph should look something like this (ignore the 80% line) – in which the reaction happens slower and produces half as much carbon dioxide gas:

In terms of collision theory, if you increase the concentration of reactants, the reaction becomes faster because it increases the frequency of collisions per second.

Changes in the Temperature of the Reaction

Do the original experiment again, but this time, at a higher temperature. Your graph will look like this (ignore the concentration label ‘cause it’s WRONG unless the lower concentration solution is still in excess).

Increasing the temperature means more kinetic energy for the particles, which make them move faster, therefore, making them collide more frequently.

Also, not all collisions make new bonds. Some particles just bounce off each other. In order for a reaction to happen, particles have to collide with a minimum amount of energy called activation energy. Increasing the temperature produces a very large increase in the number of collisions that have enough energy for a reaction to occur.

In the following diagram, a) shows a fail collision and b) shows a successful one.

Changing the Pressure

Changing the pressure of a reaction where the reactants are only solids or liquids makes virtually no difference, so the graphs remain unchanged. Increasing the pressure in a reaction where the reactants are gases does speed the reaction up. This is because it forces the particles closer together, so they hit each other more frequently.

Catalysts and How They Work

Catalysts speed up the rate of reactions but aren’t used up in the process. You can show that manganese (VI) oxide is a catalyst by simply having two conical fasks containing hydrogen peroxide. Hydrogen peroxide decomposes to give oxygen and water. Put the manganese (VI) oxide in one of the

flasks. Oxygen would be given off quickly. To check that the manganese (VI) oxide hasn’t been used up, simply filter it out from the solution and weigh it (remember to weigh it before the experiment too!). The graph should look like the pressure graph – in which the rate of reaction increases, but the amount you get at the end is still the same.

So how does it work?

Adding a catalyst gives the reaction an alternative route for reactions with a lower activation energy.

Unit 5: Quantitative Chemistry and EnergeticsA mole is a measure of the amount of substance. One mole contains 6 x 1023 (also known as the Avogadro Number) particles (atoms, molecules or formulae) of the substance. For example, 1 mol of sodium contains 6 x 1023 atoms of sodium.

Calculating Relative Atomic MassChlorine has two isotopes: chlorine-35 and chlorine-37. A typical sample will be 75% chlorine-35 and 25% chlorine-37.

The RAM = (0.75 x 35) + (0.25 x 37) = 35.5g

Calculating Relative Formula MassStraightforward stuff.

Amount in moles = Mass of Substance (g) /RFM of Element or Compound (g)

Calculations using moles:

The equation for sodium chloride is: 2Na + Cl2 2NaCl

If 2.3g of Na was used:

a) Find out how many moles of Chlorine was usedb) Find out the volume of Chlorine used in the reaction

H2 H x 2 2 x 1 = 2Ca(OH)2 (1 x Ca) + (2 x O) + (2 x

H)40 + (2 x 16) + (2 x 1) = 74

c) Find out the mass of sodium chloride produced

a) Firstly, convert the grams of Na into moles:2.3 / 23g = 0.1 mol

The equation says that 2 moles of Na and 1 mole of Cl (1 mole of a diatomic molecule is always X2) is needed to produce 2 moles of NaCl, so if 0.1 mol of Na is used, then half of that is the amount of chlorine used in the reaction in moles.

So moles of Cl used = 0.1 / 2 = 0.05 mol

b) One mole of any gas has a volume of 24 dm3 (24000cm3) at room temperature and pressure. This is also called the molar volume.

Cl2 is a gas and the moles used in the reaction = 0.05 molSo the volume of Cl2 gas used = 0.05 x 24000 = 1200cm3

c) The moles of NaCl produced is 0.1 mol (if 2 moles of Na gives 2 moles of NaCl, then 0.1 mole of Na will give 0.1 mole of NaCl). So all you do is:i) Find the RFM of NaCl (58.5)ii) Multiply that by 0.1 (5.85g)

Molar Concentrations – The Hard PartRemember that:

Mol/dm3 means moles per litre (e.g. a salt solution of 0.5 mol/dm3 means 0.5 moles (or 58.5/2 = 29.25g) of salt was dissolved in a litre of water

It’s all about proportion

20 cm3 of 0.5 mol dm3 sodium hydroxide solution was dissolved with 25cm3 of hydrochloric acid to form a sodium chloride solution. Calculate the concentration of HCl needed to react with the NaOH

NaOH + HCl H2O + NaCl

RFM of NaOH = 40g = 1 mole of NaOH0.5 mol dm3 of NaOH means (40 x 0.5) 20g of NaOH was dissolved in 1000cm3 of waterThe amount of moles in 20cm3 of NaOH solution:

20cm3/1000cm3 x 0.5 moles = 1/50 x 0.5 = 0.01 moles of NaOHThe equation says that 1 mole of NaOH + 1 mole of HCl gives 1 mole of NaClSo 0.01 moles of NaOH + 0.01 moles of HCl gives 0.01 moles of NaCl

So 0.01 moles of HCl was present in 25cm3 of HCl solution! However, concentration is measured in mol dm3 so:

1000cm3/25cm3 x 0.01 mol = 0.4 mol dm3 of HCl used.

Calculating the Empirical Formula and Molecular FormulaThe empirical formula is the simplest formula and only tells you the ratio of the various atoms. Suppose 2.4g of magnesium combined with 1.6g of oxygen, you can use a table to work out the empirical formula. (Mg = 24 O = 16)

What about with percentage figures?

Mg OCombining Masses

2.4 1.6

Number of moles 2.4/24 1.6/16= 0.10 0.10Ratio of Moles 1:1Empirical Formula

MgOC HPercentage 87.5 14.3Combining Masses

87.5 14.3

Number of moles 85.7/12 14.3/1= 7.14 14.3Ratio of Moles 1:2Empirical Formula

CH2

Suppose you had a compound containing 85.7% C, 14.3% H and you were asked to calculate the empirical formula. Firstly, you assume that 100% = 100g! (C = 12 H = 1)

However, you know that CH2 does not exist. Remember this is only the ratio. To find the molecular formulae, you need to know the relative formula mass of the compound. Suppose it was 56g for the above question.

Firstly, find out the RFM of CH2 = 12 + 2 = 14gFind out how many times 14 goes into 56, so 56/14 = 4 timesWhich means the molecular formula is C4H8!

Obtaining Formulae ExperimentallyMetal Oxides

Hydrogen can be passed over metal oxides to reduce it to the metal. To find the formula of copper oxide, the experimental steps are as follows:

1. Measure the mass of the empty combustion tube.2. Use a spatula to put copper oxide into the tube.

Weigh the tube.3. Set up the apparatus as shown. Turn the gas at

the jet to light the excess gas.4. Heat the copper oxide until it has all turned into

red copper.5. Stop heating but leave gas passing through until everything has cooled.6. Weigh the combustion tube.7. Put masses in a table and calculate empirical formula from there.

In the Case of Water of Crystallisation

When substances crystallise from a solution, water becomes chemically bounded with the salt. This is called water of crystallisation and the salt is said to be hydrated.

Suppose you had to find the formula of a BaCl2•nH2O (a barium chloride crystal), to find n:

1. Weight the mass of an empty crucible.2. Add barium chloride crystals and reweigh.3. Heat the crucible gently (so the barium chloride won’t decompose), and reweigh.4. Put masses into a table and calculate the formula from there.

Mass of Empty Tube 52.2gMass of tube + Copper Oxide (Before)

66.6g

Mass of tube + Copper 65.0Mass of Oxygen 66.6 – 65.0 = 1.6gMass of Copper 65.0 – 52.2 =

12.8g

Cu OCombining Masses

12.8 1.6

Number of moles 12.8/64 1.6/16= 0.20 0.10Ratio of Moles 2:1Empirical Formula

Cu2O

Mass of Empty Crucible 30.00gMass of tube + Crystals (Before) 32.44gMass of tube + Anhydrous Crystals (After)

32.08g

Mass of BaCl2 2.08gMass of Water 0.36g

BaCl2 H2OCombining Masses

2.08 0.36

Number of moles 2.08/208

0.36/18

= 0.01 0.02Ratio of Moles 1:2Empirical Formula

BaCl2•2H2O

Calculating Percentage YieldMost of the time, when you do carry out a chemical reaction, you get less than you expect. The rest of it has been lost in some way perhaps due to spillages or losses when chemicals are transferred.

Suppose you work out that 10g of A will give 500g of the product, but you only get 400g…?

The percentage yield is (400/500) x 100 = 80%

A general formula would be: (mass produced/expected mass to be produced) x 100

Endothermic and Exothermic Reactions

∆H represents the molar enthalpy change for exothermic and endothermic reactions

Energy CalculationsThe general formula:

Bonds of all the reactants – Bonds of all the products = Energy change

Example: Methane reacts with chlorine to produce chloromethane and hydrogen chloride. The equation:

CH4 + Cl2 CH3Cl + HCl

You would be given a table with the bonds and the energy required to break/bond them:

Reactants:

4 C – H bonds (CH4) = 4 x 413 = 1652kJ1 Cl – Cl bond (Cl2) = 1 x 243 = 243 kJ Total: 1652 + 243 = 1895 kJ

Products:

3 C – H bonds = 3 x 413 = 1236 kJ1 C – Cl bond = 1 x 346 = 346 kJ1 H – Cl (HCl) = 1 x 432 = 432 kJ Total: 2017 kJ

Endothermic Heat energy is taken inBreaking of bonds

∆H = + N kJ mol-1

Exothermic Heat energy is given outMaking of bonds

∆H = - N kJ mol-1

Bond C - H C - Cl H - Cl Cl - Cl

Energy (kJ mol -1)

413 346 432 243

(Carbon can form 4 bonds. In this case, 3 of them bonds with 3 hydrogen and the last one bonds with chlorine)

Energy Change = 1895 – 2017 = -122 kJ the reaction is exothermic

Describing Simple Calorimetry Experiments All these involve measuring a temperature change during the reaction. Specific heat is the amount of heat needed to raise the temperature of 1g of a substance by

10C. For water, the value is 4.18 J g-1 0C-1 (joules per gram per degree Celsius).

Heat Given Out = Mass x Specific Heat x Temperature Rise

For Neutralisation, Displacement Reactions and Dissolving

They all follow the same method. This example involves measuring the heat evolved (or energy) when magnesium reacts with dilute sulphuric acid.

1. Pour an excess of sulphuric acid into a polystyrene cup and measure the temperature of the acid.

2. Pour some magnesium powder into a weighing bottle and weight it.

3. Pour the powder into the acid and record the highest temperature.

4. Weigh the empty weighing bottle.

Let’s say the total mass of the solution and Mg is 50g

Heat evolved when 0.123g of Mg reacts = 50 x 4.18 x 10.1J = 2111J = 2.111kJ

To find out the heat evolved when 1 mole of Mg reacts (Mg = 24g):

(2.111/0.123g) x 24 = 412 kJ

The temperature rose, meaning the reaction is exothermic so:

Mg(s) + H2SO4(aq) MgSO4(aq) + H2(g) ∆H = -412 kJ mol-1

This is actually smaller than the accepted value, which is around -417 kJ mol-1. A reason for this could be that heat was lost too quickly. Using a mercury thermometer may give better results.

Combustion

1. Put 100cm3 of water into a conical flask and record the temperature.

2. Fill the spirit burner with alcohol (let’s say ethanol) and weight.

3. Light the spirit burner and record the temperature of water until there is say, a 400C increase.

4. Reweigh the spirit burner.

Volume of water (cm3) 100Mass of water being heated (g) 100Mass of burner before (g) 37.355

Mass of Weighing Bottle + Mg (g) 10.810Mass of Weighing Bottle Afterwards (empty) (g)

10.687

Mass of Mg used (g) 0.123Initial Temp. (0C) 17.4Final Temp. (0C) 27.5Temperature Rise (0C) 10.1

Mass of burner after (g) 36.575Mass of ethanol burnt (g) 0.780Original temp. of water (0C) 21.5Final temp. of water (0C) 62.8Water temperature increase (0C) 41.3

Heat gained = 100 x 4.18 x 41.3 = 17260 J = 17.26 kJ

Ethanol is C2H5OHOne mole of ethanol = 46g

Amount of heat produced from 1 mole of ethanol = (17.26/0.780) x 46 = 1020 kJ

KEY POINTS FOR THIS UNIT 1 mole is the Avogadro constant number of particles The molar volume is 24 dm3 or 24 000 cm3

Number of moles = mass / RFM 1 dm3 = 1000 cm3

Mol dm-3 = mol per 1000 cm3

G dm-3 = grams per 1000 cm3

Energy/Heat = mass x specific heat x temperature rise

Unit 6: Chemistry in SocietyElectrolysisRECAP! Electrolysis is a chemical change caused by passing an electric current through a compound which is either molten or in solution. An electric current (in chemistry terms), is a flow of electrons or ions. An electrolyte is a substance that undergoes electrolysis. Electrolytes all contain ions. Ionic compounds, for example, are electrolytes. Electrolytes can only undergo electrolysis when molten or in a solution, where the ions are free to move. Covalent compounds are not electrolytes because they don’t contain ions.

Electrolysis can form new substances when ionic compounds conduct electricity. It is set up as so:

The electrodes are usually made of carbon because it is fairly un-reactive.

The positive electrode is called the anode. The negative electrode is called the cathode.

A simple example of electrolysis involves molten (melted) lead (II) bromide:

So what happens?

Molten lead is found at the bottom of the cathode.

Bromine gas comes out of the anode.

When the power supply is switched off, no more bubbles are produced and everything else stops.

What the hell happened?

Since the lead (II) bromide (PbBr2) is molten, its ions are free to move around.

The bromide ions are attracted to the positive electrode. The extra electron which makes the bromide ion negatively charged is deposited into the anode, thus, turning them back into neutral bromine atoms. These then covalently bond to join bromine atoms (i.e. bromine gas).

On the other hand, the lead ions gain back to electrons (it has a 2+ charge) and become normal lead atoms. These fall to the bottom of the container as molten lead.

The half-equation at the anode would be:

2Br- Br2 + 2e-

What it basically means is that, two bromine ions are formed when one bromine molecule receives two electrons (to fill its shell). This is the format for all anions (ions that are negatively charged).

The half equation at the cathode would be:

Pb2+ + 2e- Pb

All cation half-equations are of this form. Half-equations basically show the gaining and losing of electrons. This basically means that the lead (II) ions get two electrons to become a neutral lead atom.

With molten substances, the metal will be produced at the cathode and whatever it’s bonded to will be produced at the anode.

When you electrolyse aqueous solutions (not molten salts), things are much different because you have to consider the water molecules too. Water is a weak electrolyte but it can ionise to form hydrogen and hydroxide ions.

If the metal is more reactive than hydrogen, then hydrogen ions from water is discharged instead. These pair up to form hydrogen gas that escapes as bubbles.

If the metal is below hydrogen, you get the metal produced.

If you have solutions of halides (chlorides, bromides or iodides), you get the halogen (chlorine, bromine or iodine) produced.

With other negative ions such as sulphates, oxygen would be produced.

The electrolysis of sodium chloride solution (brine) does not give sodium and

chlorine! Here’s the electrolysis:

Sodium is higher than hydrogen in the reactivity series, so hydrogen ions from the water in the sodium chloride solution is discharged instead at the cathode.

2H+ + 2e- H2

Chloride ions give up one electron each (chloride ion = 1- charge) and become chlorine atoms. These covalently bond to form chlorine gas and bubbles out of the solution at the anode.

2Cl- Cl2 + 2e-

When all the chlorine has been removed from the solution, only hydroxide (OH-) ions and sodium (Na+) ions are left, as well as some water. These combine to form sodium hydroxide solution (NaOH).

The electrolysis of sodium chloride solution is used to manufacture sodium hydroxide solution. The process is slightly different – it is electrolysed in a diaphragm cell:

The products are kept separated by the diaphragm. If the chlorine produced were to react to hydrogen, it would cause an explosion on exposure to sunlight or heat to give hydrogen chloride. Furthermore, if the chlorine were to react with the sodium hydroxide solution formed, it would form bleach. Uses of sodium hydroxide include:

Making bleach Making soap Making paper – NaOH breaks the wood down

Uses of chlorine include:

Sterilising water Making hydrochloric acid Making bleach

And… the electrolysis of copper sulphate solution:

Copper is lower than hydrogen and therefore, a coat of it forms at the cathode.

Cu2+ + 2e- Cu

Oxygen gas is discharged from the hydroxide ions in the water because the sulphate ions are more stable.

4OH- 2H2O + O2 + 4e-

If you electrolyse the solution for longer, something else happens. The hydrogen ions are being discharged and remains in the solution. Similarly, sulphate ions are being discharged either. As a result, the solution turns into sulphuric acid (H2SO4) and it begins electrolysing:

Sulphate ions are being discharged from the acid so oxygen is discharged from the hydroxide ions instead.

4OH- 2H2O + O2 + 4e-

There are only hydrogen ions arriving at the cathode so they discharge as hydrogen gas.

2H++ 2e- H2

Common Exam Question: Why is twice as much hydrogen produced than oxygen? For every four electrons that flow around the circuit, one molecule of oxygen and two molecules of hydrogen are produced.

Electrolysis CalculationsBack to moles! Here are some things to know:

One faraday means one mole of electrons passing around the circuit. One faraday = 96000 coulombs. Charge (coulombs) = Current (amps) x Time (seconds)

Example: What mass of copper is deposited on the cathode during the electrolysis of copper (II) sulphate solution if 0.15A flows for 10mins?

The electrode equation is:

Cu2+ + 2e- Cu

Calculate the coulombs involved:

10mins x 60 = 600 secondsCharge = 0.15 x 600 = 90 coulombs

The equation says that 1 mole of copper ions + 2 moles of electrons give 1 mole of copper atoms

1 mole of electrons = 96000 coulombs2 moles of electrons = 192 000 coulombs2 moles of electrons (192 000 coulombs) give 1 mole of copper (RFM = 64g), so 90 coulombs give:

(90/192 000) x 64g = 0.03g

(Coulombs worked out/coulombs of electrons) x RFM of element = Mass of element deposited

When involving gases…

Example: During the electrolysis of dilute sulphuric acid, hydrogen is released at the cathode and oxygen at the anode. Calculate the volumes of hydrogen and oxygen produced if 1.0A flows for 20mins

The electrode equations are:

2H++ 2e- H24OH- 2H2O + O2 + 4e-

Assume the molar volume of gas to be 24000 cm3

For hydrogen:

2H++ 2e- H2

2 hydrogen ions + 2 moles of electrons give 1 mole of hydrogen molecule

20mins x 60 = 1200 seconds1200 x 1A = 1200 coulombs

2 x 96 000 coulombs (2 moles of electrons) = 192 000 coulombs1 mole of hydrogen gas = 24000 cm3

192 000 coulombs give 1 mole of hydrogen gas (or 24000 cm3 of hydrogen gas)So 1200 coulombs give: (1200/192 000) x 24000 = 150 cm3 of hydrogen produced

For oxygen:

4OH- 2H2O + O2 + 4e-

20mins x 60 = 1200 seconds1200 x 1A = 1200 coulombs

4 moles of hydroxide ions give 2 moles of water + 1 mole of oxygen gas + 4 moles of electrons (4 x 96000 = 384 000 coulombs)

1 mole of oxygen gas = 24000 cm3

384 000 coulombs give 1 mole of oxygen gas (or 24000 cm3 of oxygen gas)So 1200 coulombs give: (1200/384 000) x 24000 = 75 cm3 of oxygen produced

The equation:

(Calculated coulombs/moles of electrons) x (mole of gas x 24000) = amount of gas produced in cm3

Reversible Reactions and Dynamic EquilibriaSome reactions are reversible. Reversible reactions are indicated by the symbol

Some examples of reversible reactions include:

Copper (II) Sulphate Crystals

Heating the blue hydrated copper (II) sulphate crystals causes them to lose their water of crystallisation, making them turn from blue to white – the white copper (II) sulphate crystals are described as ‘anhydrous’ meaning ‘without water’:

CuSO4•5H2O CuSO4 + 5H2O

However, this reaction can be reversed by simply adding water to the crystals. The crystals will become hydrated again:

CuSO4 + 5H2O CuSO4•5H2O

Heating Ammonium Chloride

When ammonium chloride is heated in a test tube, the white crystals decompose into hydrogen chloride gas and ammonia gas. These flow upwards and recombine again further up the test tube:

NH4Cl HCl + NH3

This later recombines:

HCl + NH3 NH4Cl

Introducing Dynamic EquilibriaThings change when reversible reactions are carried out under ‘closed’ conditions – meaning no substances are added to the reaction mixture and no substances can escape from it. Heat however, can be given off or absorbed.

In a reversible reaction, you have the forward reaction (the reaction going from left to right) and the back reaction (the opposite of the forward reaction) happening at the same time. Both rates of reactions will become equal and this point is the dynamic equilibrium. It is dynamic in a sense that the reactions are still continuing, and equilibrium in a sense that the total amounts of the various things present are now constant. In other words:

A + 2B C + D

When you have a reaction like the above, A + 2B (forward reaction) is reacting to produce C + D (back reaction). At the same time, C + D is reacting to produce A +2B. In the end, you have equal amounts of products and reactants. Another way to think of is, is to imagining walking down an elevator that

goes up, making sure you’re walking at the same speed as the elevator. You would be going down, but everytime you take one step down, the elevator goes one step up. In the end, you remain where you are.So how would you produce more of substance C in a reversible reaction such as the above? You can do this by altering the position of the equilibrium by either:

Changing the pressure Changing the temperature Increasing/Decreasing the concentrations of substances present Adding a catalyst

If a dynamic equilibrium is disturbed by changing the conditions, the reaction moves to counteract the change.

In other words, the reaction will either go more towards the ‘forward’ direction or the ‘back’ direction in an attempt to ‘adapt’ to the conditions.

A + 2B C + D

Changing the Concentration

What happens when more of A is added? If you add more A, the reaction will want to remove it. This can only be done by reacting more A to 2B, and in the end, this gives us more C and D. The conditions, in this case, favour the forward reaction.

Changing the Pressure

When you increase the pressure, you bring molecules closer together. Increasing the pressure will always help the reaction go in the direction which produces the smaller number of molecules. In this case, we have 3 molecules on the left (one A and 2 Bs), whereas, we only have 2 molecules on the right (one C and one D). The reaction can only reduce the pressure by producing few molecules. This can only be done by producing more C and D.

Changing the Temperature

Suppose the forward reaction was exothermic

A + 2B C + D ∆H = -100 kJ mol-1

This would mean that the back reaction would be endothermic by the same amount (-100 kJ mol-1).

Suppose the temperature was decreased, the reaction would respond by increasing the temperature back up again. This can only be done by producing more C and D because the forward reaction is exothermic. Increasing the temperature will of course, have the opposite effect.

Adding a Catalyst

Adding a catalyst speeds up the forward and back reactions by the same proportion. This means that there is no change in the position of the equilibrium.

Methods of ExtractionThe extraction of metals depend alot on its position in the reactivity series. Costs are also factors to take into account. For metals up to zinc, the cheapest method is usually heating the ore with carbon or carbon monoxide to reduce it. For metals more reactive than zinc, electrolysis is usually used.

The Extraction of AluminiumAluminium is extracted from an ore called bauxite, which is impure aluminium oxide. Aluminium ions are attracted to the cathode and are reduced to aluminium:

Al3+ + 3e- Al

Oxide ions are attracted to the anode and lose electrons to form oxygen gas:

2O2- O2 + 4e-

There are a few things to keep in mind:

Melting aluminium oxide requires extremely high temperatures. Instead, it is dissolved in an aluminium compound called cryolite.

Because of high temperatures, the carbon anodes will burn with oxygen to form carbon dioxide. This means the anodes have to be replaced regularly.

The cost of electricity is also a major factor – the cell has currents up to 100 000A so it is expensive.

Uses of aluminium include:

Use Property of Aluminium That Makes This UsefulTo make aircraft carriers Resists corrosion due to its aluminium oxide coat

Has low density Strong

Saucepans Aluminium has a shiny appearance It is a good conductor of heat It resists corrosion It has a low density

The Iron Blast FurnaceIron is extracted from an ore that contains iron (III) oxide called haematite.

Coke is impure carbon. It burns to form carbon dioxide. This is a strongly exothermic reaction.

C + O2 CO2

At high temperatures, the carbon dioxide is reduced by more carbon to give carbon monoxide.

C + CO2 2CO

Carbon monoxide is the main reducing agent:

Fe2O3 + 3CO 2Fe + 3CO2

Carbon may also reduce the iron (III) oxide:

Fe2O3 + 3C 2Fe + 3CO

The heat of the furnace causes the limestone to thermally decompose to form calcium oxide and carbon dioxide:

CaCO3 CaO + CO2

The calcium oxide reacts with silicon dioxide (one of the impurities found in haematite) to form calcium silicate, which melts and trickles to the bottom of the furnace as molten slag:

SiO2 + CaO CaSiO3

Uses of Iron (you don’t really need to learn this)Types of Iron Iron Mixed With Some UsesWrought Iron Pure Iron Decorative work such as gates and

railingsMild Steel 0.25% Carbon Nails, car bodies, ship buildingHigh – Carbon Steel 0.25 – 1.5% Carbon Cutting toolsCast Iron About 4% Carbon Manhole covers, guttering, engine

blocksStainless Steel Chromium and Nickel Cutlery, cooking utensils

Preventing the Rusting of Iron (in which iron oxidises into iron oxide Fe2O3)Using Barriers Keep water/oxygen away from the iron by painting, coating with oil…etc.Alloying the Iron Such as allowing it with chromium and nickel to produce stainless steelUsing Sacrificial Anodes

Galvanising iron by coating it with a layer of zinc. Zinc is more reactive than iron and will corrode instead. During the process it loses electrons to form ions. These electrons flow into the iron so any iron atom which has lost electrons immediately regains them. These means even if the zinc is scratched, the iron won’t rust.

The Haber ProcessThe Haber process is used to make ammonia NH3

Uses of ammonia include:

Making fertilisers Making nitric acid Making nylon

The equation for the reaction is:

N2 + 3H2 2NH3 ∆H = -92 kJ mol-1

The forward reaction would be favoured by a low temperature because the forward reaction is exothermic (so lowering the temperature would cause the reaction to make more NH3 to heat things up abit more). 450ºC isn’t a low temperature. It is however, a compromise temperature, because if the temperature was made to be low, the reaction would be so slow that it would take a very long time to produce much ammonia.

Pressure is also another compromised. Because the forward reaction has less molecules than the back reaction (2 molecules of NH3 as opposed to 1 N2 and 3 H2 molecules), the forward

reaction would be favoured by a high pressure. 200 atm is high, but anything higher would be extremely expensive.

The iron catalyst speeds the reaction up but has no effect on the equilibrium. However, if the catalyst wasn’t used, the reaction would be too slow.

The Contact Process1. Burn sulphur in air to form sulphur dioxide SO2

S + O2 SO2

2. Use an excess of air to react sulphur dioxide to more oxygen to form sulphur trioxide

2SO2 + O2 2SO3 ∆H = -196 kJ mol-1

3. Reacting sulphur trioxide with water will give an uncontrollable fog of concentrated sulphuric acid. Instead, sulphur trioxide is absorbed in concentrated sulphuric acid to give fuming sulphuric acid (oleum):

SO3 + H2SO4 H2S2O7

4. This is converted into twice as much concentrated sulphuric acid by careful addition of water:

H2S2O7 + H2O 2H2SO4

The reversible reaction here is:

2SO2 + O2 2SO3 ∆H = -196 kJ mol-1

Because the forward reaction is exothermic, a low temperature has to be used. Again, if a low temperature is used, the rate of reaction would be too slow, so 450ºC is a compromise.

As for the pressure, a low pressure is needed because the forward reaction contains fewer molecules than the back reaction.

The catalyst, vanadium (V) oxide (V2O5) speeds up the rate of reaction but has no effect on the equilibrium. Again, without a catalyst, the rate of reaction would be extremely slow.

Uses of sulphuric acid include:

Making fertilisers – including ammonium sulphate and other substances

Detergents – including hand soaps and shampoos

Paint Manufacture – used to extract titanium dioxide from titanium ores

A Few Extra Bits and PiecesIndicator SolutionsNote: Methyl orange is orange in neutral solutions, however, these two indicator solutions are both yes no indicators – meaning, if there is a reaction, the methyl orange (for example) would turn from red to yellow (or vice versa depending on whether the solution it has been dropped into is acidic or alkaline). It will not turn orange.

The acidity or alkalinity of something is measured in pH. Universal indicator cam be used to measure the approximate pH of a solution.

Acids are sources of H+ ions Alkalis are sources of OH- ions for more info, go to page 10

Reactions Between Metals, Metal Compounds and AcidsBasic things to know:

Hydrochloric acid reacts with metals and metal compounds to form a metal chloride Sulphuric acid reacts with metals and metal compounds to form a metal sulphate

And some equations:

Metal + Acid Salt + Hydrogen

Mg + 2HCl MgCl2 + H2Mg + H2SO4 MgSO4 + H2

Metal Oxide + Acid Salt + Water

MgO + 2HCl MgCl2 + H2OMgO + H2SO4 MgSO4 + H2O

Metal Carbonate + Acid Salt + Water + Carbon Dioxide

MgCO3 + 2HCl MgCl2 + H2O + CO2Mg CO3 + H2SO4 MgSO4 + H2O + CO2

Preparing Soluble Salts Using TitrationThe exact amount of acid needed to neutralise an alkali can be foundby titration. This technique can be used to make pure crystals of a soluble salt (one that dissolves in water).

Phenolphthalein colourless in acid

pink in alkali

Methyl orange red in acid yellow in alkali

In the example below, an acid and an alkali react to make sodium chloride.

1. The burette is filled with hydrochloric acid.

2. A known quantity of alkali (say 50 cm3 sodium hydroxide) is released from a pipette into the conical flask. The tap on the burette is turned open to allow the acid to be added drop by drop into the alkali.

3. The alkali contains an indicator (phenolphthalein).

4. When enough acid has been added to neutralise the alkali, the indicator changes from pink to colourless. This is the end point of the titration.

5. The titration can be repeated using the same amounts of acid and alkali but without the indicator.

6. Pure salt crystals which are free from indicator can then be crystallised from the neutral solution.

Precipitation ReactionsThe process of making a solid come from a solution is called precipitation. The solid itself is called a precipitate. An insoluble salt (one that doesn't dissolve) can be made by reacting the appropriate soluble salt with an acid or alkali or another salt.

You are normally asked to prepare a solid from two soluble solutions – so know your solubility rules:

All nitrates are soluble. All sodium, potassium and ammonium compounds are

soluble. Most carbonates and hydroxides are insoluble except

for sodium, potassium and ammonium. All sulphates are soluble except barium and lead(II)

sulphate. All chlorides are soluble except lead(II) and silver

chloride.

Example: Prepare Silver Chloride

You’ll need:

A soluble silver salt – what about silver nitrate? A soluble chloride – like magnesium chloride?

You can make it up really. So the equation:

AgNO3 (aq) + MgCl2(aq) AgCl(s) + Mg(NO3)2(aq)The Reactivity SeriesMetals are arranged based on their reactions in the reactivity series.

They can be deduced by using displacement reactions, in which a less reactive metal is pushed out of its compound by a more reactive metal. For example, the reaction between magnesium and copper (II) oxide:

Mg + CuO MgO + Cu

Displacement reactions are examples of redox reactions, in which oxidation and reduction occurs in the same reaction.

STATE SYMBOLS(s) solid(l) liquid(g) gas(aq) aqueous solution

(dissolved in water)

The reducing agent is a substance that reduces something else. In this case, it’s magnesium. The oxidising agent is a substance that oxidises something else. In this case, it’s copper.

Remember OILRIG: Oxidation is gain (of electrons); Reduction is loss (of electrons)

Let’s look at this in terms of an ionic equation:

Mg + Cu2+ + O2- Mg2+ + O2- + Cu

This basically looks at what turns into ions and what does not. The magnesium turns into a positive ion because it loses two electrons to bond with oxygen. The copper ion, because it’s displaced, regains the two electrons it lost to the oxygen.

Notice in this equation, the oxygen ion does not change? It remains an ion. In this case, the oxygen ion is the spectator ion. Spectator ions aren’t included in the ionic equation, so the proper ionic equation of this displacement reaction should be:

Mg + Cu2+ Mg2+ + Cu

As mentioned earlier, the magnesium loses two electrons to bond with oxygen and copper gains two electrons. This can be written as half-equations:

Mg Mg2+ + 2e-

Cu2+ + 2e- Cu

Reactions with substances are different depending on its reactivity:

Metals Reaction with Water Reaction with Dilute AcidsPotassium, Sodium, Lithium

Very vigorous, produces hydroxides and hydrogen gas

Too reactive to add safely to acids

Calcium Reacts gently and produces the same products as above

Can be added to very dilute acids – but it’s going to be violent!

Magnesium Reacts with steam to produce magnesium oxide and hydrogen

Reacts vigorously – strongly exothermic – forms a salt and hydrogen

Zinc or Iron Reacts slowly and forms an oxide and hydrogen

Reacts slowly

Anything Below No reaction No reaction

Rusting of IronThis requires two things:

Oxygen Water

Preventing the Rusting of Iron (in which iron oxidises into iron oxide Fe2O3)Using Barriers Keep water/oxygen away from the iron by painting, coating with oil…etc.Alloying the Iron Such as allowing it with chromium and nickel to produce stainless steelUsing Sacrificial Anodes

Galvanising iron by coating it with a layer of zinc. Zinc is more reactive than iron and will corrode instead. During the process it loses electrons to form ions. These electrons flow into the iron so any iron atom which has lost electrons immediately regains them. These means even if the zinc is scratched, the iron won’t rust.

More on sacrificial anodes – if the iron has already rusted, it can still be displaced by the more reactive zinc:

Fe2O3 + 3Zn 2Fe + 3ZnO

And this is the end.

Good luck!