Products from crude oil

1. Chemicals from oil

How oil is formed

Oil is thought to have formed over millions of years from the break down of tiny dead creatures. Natural gas is formed alongside oil.

The dead organisms sank to the bottom of lakes or seas and became trapped in muddy sediments. As the sediments built up, the lower layers were under pressure. They eventually turned to rock. If there was no oxygen in the sediments, heat and pressure turned the remains of the organisms into oil and natural gas.

Some rocks are porous - they have a network of tiny holes in them.Sandstone and limestone are examples. Oil is a liquid so it seeps into porous rocks. Gas also diffuses into these rocks.

Porous rocks may also contain water. Gas and oil do not mix with water. They are less dense than water. This means they form layers above the water.

Sometimes the rock layers form so that the oil and gas are trapped under the rock such as shale that is not porous. Large amounts of oil and gas may collect in a porous rock. The pressure on the oil may build up so much that when a hole is drilled through the rock cap, oil gushes out.

Fractional distillation of crude oil

Crude oil is a mixture of many thousands of different compounds with different properties. They are called hydrocarbons because they only contain the elements hydrogen and carbon.

To make crude oil useful, batches of similar compounds with similar properties need to be sorted. These batches are called fractions and they are separated by fractional distillation.

The theory behind this technique is that some of the compounds in crude oil are easily vaporised, for example, they are volatile due to their low boiling points. Others are less volatile and have higher boiling points.

In fractional distillation, the crude oil is heated to make it vaporise. The vapour is then cooled. Different fractions of the oil are collected at different temperatures.

Fraction:	No. of carbon atoms:	Colour:	Boiling point range oC:	Uses:
Refinery gas	1 - 4	Colourless	Below room temp.	Gaseous fuel, making chemicals.
Gasoline (petrol)	4 - 12	Colourless to pale yellow	32-160oC	Motor car fuel, making chemicals.
Kerosine (paraffin)	11 - 15	Colourless to yellow	160-250oC	Heating fuel, jet fuel.
Diesel oil	15 - 19	Brown	220-350oC	Diesel fuel for lorries, trains, etc. and heating fuel.
Residue lubricating oil heavy fuel oil bitumen	C	Dark brown	Above 350oC	Fuels for power stations, ships etc. Some is distilled further to give lubricating oils, waxes, etc.
	20 - 30
	30 - 40
	50 and above

As the hydrocarbon molecule chain increases its boiling point increases, it becomes more viscous, becomes more difficult to light, the flame becomes sootier and it develops a stronger smell.

2.Products from crude oil

Alkanes

Physical properties:

The chemistry of carbon compounds is called organic chemistry. There are millions of organic chemicals, but they can be divided into groups called homologous series. All members of a particular series will have similar chemical properties and can be represented by a general formula.

The alkane series is the simplest homologous series. The main source of alkanes is from crude oil.

Alkanes are covalent compounds. They are hydrocarbons, which means they contain hydrogen and carbon. The general formula for an alkane is C_nH_2n+2.

Properties and uses of alkanes:

Name of alkane:	Melting point oC:	Boiling point oC:	Density g/cm3:	State at room temperature:
Methane CH4	-182	-162	0.42	Gas
Ethane C2H6	-183	-88	0.55	Gas
Propane C3H8	-188	-42	0.58	Gas
Octane C8H18	-57	126	0.72	Liquid

The first four alkanes are gases at room temperature.

Alkanes with 5-17 carbon atoms are liquids.

Alkanes with 18 or more carbon atoms are solids.

As the number of carbon atoms increases, the melting points, boiling points and densities increases.

They are insoluble in water but dissolve in organic solvents such as benzene.

Their chemical reactivity is poor. The C-C bond and C-H bond are very strong so alkanes are not very reactive.

They will carry out combustion. Burning alkanes in air (oxygen) produces water and carbon dioxide. The reactions are very exothermic (give out heat energy), so alkanes in crude oil and natural gas are widely used as heating fuels.

For example:

If alkanes combust in too little air, carbon monoxide may form. This is dangerous and can cause death.

Cracking alkanes

The lighter fractions (for example, petrol) are in large demand. The heavier fractions are not so useful but unfortunately chemists have to be able to convert these heavier fractions into petrol and other useful products, due to supply and demand, by a method known as cracking.

Cracking breaks down molecules into smaller ones. Catalysts or heat may be used to crack the alkane chain into smaller ones.

Note, that one of the products that is formed when we crack naphtha contains a double bond between two carbon atoms. A hydrocarbon that possesses one double bond belongs to the next homologous series called alkenes.

Another reaction that often occurs after fractional distillation is reforming. Hydrocarbons of the same formula have different boiling points. Straight-chained alkanes have greater boiling points than the branched version. This means they catch light more easily - but this can be too much for the hot cylinder of the car engine. Reforming converts straight-chained alkanes to branched.

Alkenes

The members of this series contain a double bond. They are hydrocarbons.

The general formula of the alkenes is C_nH_2n Most alkenes are formed when fractions from the fractional distillation of crude oil are cracked.

Properties of alkenes:

Like alkanes, the boiling point, melting point and densities increase with larger size molecules.

They are insoluble in water.

They combust like alkanes to produce carbon dioxide and water. However, they burn with sootier flames due to their higher percentage of carbon content to hydrogen.

Chemically, alkenes are more reactive than alkanes. This is because they possess a double bond that can be broken open and added to in a reaction.

For example:

These reactions are called addition reactions.

Saturated and unsaturated:

Organic compounds, like alkanes, which have four single covalent bonds to all their carbon atoms are described as saturated.

Alkenes are hydrocarbons with a double bond between two carbon atoms and are described as unsaturated. This is because they do not have the maximum number of atoms attached to their four bonds, as one is double!

Polyunsaturated margarines and vegetable oils contain many C=C bonds.

3. Polymerisation

Making plastics

Facts about plastics:

Polythene (polyethene) is made by forming a long chain of ethene molecules. Many other compounds are made in a similar way. A compound made like this is called a polymer.

Polymers are long chains of monomers. A monomer is the building block or in other words the repeating unit that is used to make the polymer. In the above example, ethene is the monomer and polythene the polymer.

Polystyrene (many styrene molecules) is another well-known polymer.

Many polymers can be easily moulded into many shapes - these are called plastics.

Polymerisation is the name given to the reaction that produces polymers.

Remember: alkenes can become polymers but alkanes cannot. This is because alkanes are saturated whereas alkenes are unsaturated which means that they can carry out addition reactions, required for polymerisation.

This type of polymerisation is called addition polymerisation.

ACIDS AND ALKALIS

1.Acids, Alkalis and Neutral Substances

Properties of acids

1. They are liquids.

2. They are solutions of compounds in water.

3. If concentrated they can be corrosive.

4. Acids taste sour (for example, vinegar).

5. Turn blue litmus paper red - this is an easy test for an acid!

6. Usually react with metals to form salts.

7. Acids contain hydrogen ions.

8. Turn Universal Indicator from green to red, and have a pH less than 7.

Examples of acids: are vinegar (ethanoic acid) and lemon juice (citric acid)

magnesium + hydrochloric acid -> magnesium chloride + hydrogen gas

Some common acids used in your laboratories at school will be:

1. Hydrochloric acid, HCl_(aq)

2. Nitric acid, HNO_3(aq)

3. Sulphuric acid, H₂SO_4(aq)

Properties of alkalis

1. They feel soapy to touch.

2. They are soluble bases.

3. Like acids, they can burn the skin.

4. They turn red litmus blue - this is how you test for an alkali!

5. Alkalis contain hydroxide ions (OH^-).

6. They taste bitter.

7. Turns Universal Indicator from green to blue or purple.

Some common alkalis used in your laboratories at school will be:

1. Sodium hydroxide, NaOH_(aq)

2. Ammonia, NH₃NH₄OH_(aq)

3. Calcium hydroxide, Ca(OH)_2(aq)

Properties of neutral substances

1. Litmus paper is not affected by neutral paper.

2. Tend to be harmless

3. Universal Indicator stays green.

Common examples of neutral substances:

1. Water

2. Sodium chloride solution, NaCl_(aq)(common salt)

3. Sugar solution C₆H₁₂O_6(aq)

The pH scale

The Strength of an Acid

Acids and alkalis can be strong or weak!

So how can we measure their strength?

The strength of an acid or alkali is shown using a scale of numbers called the pH scale. The numbers go from 0-14.

On the scale it follows that:

An acidic solution has a pH number less than 7

An alkaline solution has a pH number greater than 7

A neutral solution has a pH number of exactly 7.

You can find the pH of any solution using universal indicator. Universal indicator is a mixture of dyes. It comes as a solution or in paper.

Universal indicator will change from green to a different colour depending on the pH of the solution you place it in.

Note:

In a strong acid, nearly all the acid molecules form ions.

In a weak acid, only some of the acid molecules form ions.

The more OH^- ions (hydroxide ions), the more alkaline an alkali will be.

In other words, the more OH- ions there are the higher the pH number.

2. Neutralisation

Reactions between acids and alkalis

The Making of a Salt

When an acid reacts with an alkali it produces a salt and water.

This reaction is called neutralisation. The alkali has neutralised the acid by removing its H+ ions, and turning them into water.

Neutralisation always produces a Salt

Uses of Neutralisation

Soil Treatment - Farming

The majority of plants grow best at pH 7. If the soil is acidic or alkaline the plant may grow badly. Therefore, chemicals can be added to the soil to change its pH.

If the soil is too acidic - the most common complaint - it is treated with a base (chemicals opposite to an acid) in order to neutralise it. Common treatments use quicklime (calcium oxide) or chalk (calcium carbonate).

Indigestion

We all have hydrochloric acid in our stomach - it helps breakdown food! However, too much acid leads to indigestion. Therefore, to cure this ailment we need to neutralise the acid with a base such as, sodium hydrogen carbonate (baking soda), or an indigestion tablet.

Insect Stings

A bee sting contains acid. In order to relieve the painful symptoms of the sting we need to neutralise the acid. By rubbing on calamine lotion (zinc carbonate) or baking soda the acid can be neutralised.

Wasp stings are alkaline, hence acid is needed to neutralise and remove the painful sting. Vinegar (ethanoic acid) is used.

Waste from Factories

Waste from many factories are often acidic. If this acidic solution is not treated and enters rivers it can kill fish. Slaked lime (calcium hydroxide) is often used to neutralise the acid.

Chemical Bonding test

MULTIPLE CHOISE

1.What happens when a covalent bond is formed?

A. Electrons are transferred from one atom to another

B. Electrons are shared between two atoms

C.Protons are transferred from one atom to another

D.Protons are shared between two atoms

2.When magnesium makes an ionic bond with oxygen it loses two electrons. What is the charge on the magnesium ion?

a. -2

b. 0

c. +2

d. +1

3.What type of bond is found in ammonia, NH₃?

a. ionic

b. hidrogen

c. covalen

d. dative

4. What are all atoms trying to achieve when they bond with other atoms?

a. To lose all of their protons

b.To lose all of their electrons

c.To have a full shell of protons

d.To have a full shell of electrons

5. How many electrons are shared between the two atoms of a hydrogen molecule, H₂O (Atomic number=1)

a. 1

b. 2

c. 3

d.4

6. How many electrons are shared between the two atoms of a nitrogen molecule? (Atomic number = 7)

a. 2

b. 4

c. 6

d. 8

7. What kind of structure is found in diamond?

a. Giant molecular

b. simple molecular

c. ionic

d. covalent

8. Graphite the only non-metal element that is able to conduct electricity. How can it do this?

a. It has a regular arrangement of 4 covalent bonds from one atom of carbon to 4 other atoms

b. It has free electrons floating between sheets of covalently bonded carbon atoms

c. It has free electrons floating around fixed positive carbon ions

d. It is made up of ions that are able to move through the lattice

9. Which of the following descriptions of metallic bonding is true?

a. Metallic bonding consists of a sea of electrons floating around fixed cations

b. Metallic bonding consists of a sea of electrons floating around fixed anions

c. Metallic bonding consists of thousands of atoms held together by shared pairs of electrons

d. Metallic bonding consists of lots of ions being held together in a lattice

10. The melting point for a compound is 1570^oC. Which type of bonding is present in the compound?

a. ionic

b. covalent

c. hidrogen

d. dative

Chemical Bonding 2

COVALENT BONDING

Sharing electrons

When two non-metals react together, they both need to gain electrons to complete full outer shells. The only way this can be achieved is if they share their outer electrons

Hydrogen: Each hydrogen atom has only one electron and needs one more to complete its first shell. When two hydrogen atoms get close together their shells can overlap and then they can share their electrons.

Since, electrons are being shared, there is a strong force of attraction between them. This force is a covalent bond.

The bonded atoms form molecules. Hydrogen's molecular formula is H₂.

Chlorine: A chlorine atom needs a share of one other electron to obtain a full outer shell. If two chlorine atoms are placed together the result is as shown below:

Oxygen: Each oxygen atom requires a share of two electrons.

Since each oxygen atom has a share of two pairs of electrons, we call this a double covalent bond.

Covalent compounds

There is a vast number of compounds that exist as molecules.

Water: In each molecule, H₂O, one oxygen atom shares electrons with two hydrogen atoms.

Ammonia: In each molecule, NH₃, one nitrogen atom shares its electrons with three hydrogen atoms, so that they all reach full shells.

Methane: Its formula is CH₄. One carbon atom shares its electrons with four hydrogen atoms.

2. TYPES OF SOLID

The four types that we shall study in this quick learn are:

1. Metals

2. Ionic

3. Molecular

4. Giant molecular

The metals

In a metal, the atoms are very tightly packed, leaving little space between them. Due to this tight packing, the outer electrons become delocalised from their atoms. This results in a 'sea' of electrons around a lattice of ions or 'pseudo' cations.

Properties of metals

Here are some general properties, but remember there are always exceptions!

1. They are hard.
2. They are tough.
3. They are not easily compressed.
4. High tensile strength - not easily stretched.
5. Malleable - can be bent or hammered into a shape.
6. Ductile - can be drawn into wires.
7. Good conductors of heat and electricity because of sea of electrons that can move around the lattice carrying heat energy or charge.
8. Usually high melting points.

Ionic solids

Ionic solids are made up of a lattice composed of oppositely charged ions. One of the most common ionic solids is sodium chloride. Sodium chloride is made up of sodium and chloride ions packed - a lattice. The ions are held by electrostatic charges in an ionic bond.

Properties of ionic solids

1. High melting points and boiling points due to strong ionic bonds. Most are solids at room temp.
2. They are brittle - will shatter with a hammer.
3. Usually soluble in water. Insoluble in non-polar solvents.
4. Do not conduct electricity in solid state. They do conduct when molten or dissolved in water since the ions are free to carry the charges as the ionic bonds do not hold them firmly in the liquid state.

Molecular solids

In a molecular solid, the molecules are held together by weak Van Der Waal's force, but packed in a regular pattern. Iodine is an example of a molecular solid. Each iodine molecule is made up of 2 iodine atoms, held together by a strong covalent bond. Each iodine molecule is held to another by weak Van Der Waal's forces.

Properties of molecular solids

1. Low melting and boiling point due to weak forces between molecules.
2. They are brittle.
3. Insoluble in water but soluble in non-polar solvents such as tetrachloromethane and petrol.
4. Do not conduct electricity. Molecules do not carry a charge so even when melted, molecular solids cannot conduct.

Giant molecular solids

Diamond and graphite

Diamond: Is made up of a lattice of carbon atoms. Each carbon atom can make 4 covalent bonds to 4 other carbon atoms. Each outer atom then bonds to 3 more and so on. Eventually millions of carbon atoms are bonded to form a giant lattice.

Properties of diamond

1. Very hard - hardest known substance. Each atom held to 4 others by strong covalent bonds - this explains the high melting point.
2. Does not conduct electricity due to no ions or free electrons to carry charge.

Graphite: Graphite is made up of flat sheets of carbon atoms.

Each carbon atom makes three covalent bonds to other carbon atoms. T his gives rings of 6 atoms. The flat sheets that lie on top of each other are held by weak forces - Van Der Waal's.

Properties of graphite

1. It is soft and slippery due to sheets of atoms been able to slide over one another because of weak forces between them.
2. A good conductor of electricity. This is due to each atom only using 3 out of 4 outer electrons in bonding. The fourth electron of each atom becomes delocalised throughout the lattice, enabling graphite to carry charge.
3. High melting point due to strong covalent bonds holding atoms of carbon together in the rings.

Chemical Bonding 1

How are compounds formed?

formation of compounds

Most elements form compounds.

For example: A reaction between sodium and chlorine gives the compound sodium chloride (salt) quite readily.

The noble gases do not usually form compounds. They are different from other elements, since their atoms are described as stable or unreactive. They are stable because their outer electron shell is full. A full outer shell makes an atom more stable.

Only the noble gases have full outer shells. This is why they are stable.

Other elements react with each other in order to obtain full outer shells, this makes them more stable.

How atoms lose and gain electrons

Depending on their electronic configurations, atoms lose or gain electrons in order to achieve a full outer shell.

Losing electrons

The sodium atom has one electron in its outer shell. If it loses this one electron it will achieve a full outer shell. By losing the one electron to another atom, it becomes a sodium ion.

The sodium ion still has 11 protons but by losing one electron it has only 10 electrons compared to the atom. Hence, its overall charge is +1.

This +1 charge is due to the ion having one more proton than electron.

In naming ions, you take the symbol Na and assign a positive charge. This gives us the sodium ion Na⁺.

Gaining electrons

A chlorine atom has seven electrons in its outer shell. It can reach a full outer shell by gaining one electron. It will then become the chloride ion, Cl-.

A negative charge is assigned to the ion to signify that the ion contains one more electron than proton.

Ions

Any atom can become an ion if it gains or loses electrons.

An ion is a charged particle. It is charged due to an unequal number of electrons and protons.

2. IONIC BONDING

Remembering that elements gain or lose electrons, when forming compounds,so that they achieve a full outer shell - let's now look at the reaction between sodium and chlorine.

Example 1: Reaction between sodium and chlorine

A sodium atom loses one electron to achieve a full outer shell and chlorine gains one electron to complete a full outer shell. So when a sodium atom reacts with a chlorine atom, the sodium atom loses its one electron to chlorine. The two ions formed are a sodium ion, Na⁺ and a chlorine ion Cl^-.

The two ions have opposite charges, they attract one another.

The force of attraction between them is an electrostatic one. This type of attraction is strong. It is called an ionic or electrovalent bond.

Example 2: Reaction between magnesium and oxygen

Other metals and non-metals react together to form ionic compounds This is because metals tend to lose electrons, whereas non-metals tend to gain electrons.

A magnesium atom has two electrons in its outer shell, whereas oxygen has six electrons. This means that magnesium wants to lose two (to oxygen) and oxygen wants to gain two (from magnesium) so that they can have full outer shells.

The ions attract each other due to their opposite charges. Magnesium ions and oxide ions are formed. The product is magnesium oxide, MgO.

Example 3: Reaction between magnesium and chlorine

To obtain full outer shells magnesium must lose two electrons and chlorine must gain one electron. So when we react magnesium in chlorine, one magnesium atom reacts with two chlorine atoms to form magnesium chloride, MgCl₂.

The Periodic Table

Group I - the alkali metals

Lithium, sodium and potassium all belong to Group 1.

This is because they all have 1 electron in their outer shell which is why they react in similar ways.

Properties:

Soft metals that can be cut with a knife.

Low density - can float on water.

Low melting points in comparison with other metals.

They react violently (in some cases) with water to form alkaline solutions - hence the name, alkali metals.

Reactivity increases as you descend the group.

Potassium is more reactive than lithium, since although they both need to lose one electron to have full outer shells, potassium's outer electron is furthest from the positive attractions of the nucleus. Therefore, it is easier for potassium to lose its outer electron than it is for lithium.

Other trends:

Melting point and boiling point decreases down the group.

Group II - the alkaline earth metals

Magnesium, Calcium and Strontium all belong to Group 2.

All Group 2 elements have two outer electrons, therefore they wish to lose two when bonding to create compounds. Losing two electrons allows them to have full outer shells, and achieve stability.

Properties:

Silvery metals.

Higher melting and boiling points than Group I elements.

Less reactive than Group I elements. This is because it is more difficult to lose two electrons compared to losing just one electron.

React with water to form alkaline solutions. Reactivity increases down the group. This is because the smaller the atom the closer the outer electrons are to the nucleus. Therefore there is a greater attraction between the nucleus and electrons in magnesium than there is in calcium.

Melting points and boiling points decrease down the group due to weaker forces of attraction between atoms.

Hardness increases as you descend down the group.

The Periodic Table

Placing elements in order

How the Periodic Table was designed

Scientists have managed to place all the elements in order using the following system:

They are placed in order of increasing proton number.

Hydrogen is first with a proton number of 1.

Next the list of elements are picked out and placed into groups. Groups are defined by the number of electrons in the outer shell of the atom.

Example: Lithium has one electron in its outer shell so is placed into Group 1. Atoms with two outer shell electrons are placed in Group 2. There are eight groups in total.

Finally, the groups are placed next to one another to finally complete the Periodic Table:

The groups

As has already been stated there are eight groups in the Periodic Table plus, a block of elements called the Transition Metals.

Some of the groups have special names:

Group 1: the alkali metals.

Group 2: the alkaline earth metals.

Group 7: the halogens.

Group 0 or 8: the noble gases.

The zigzag line through the Periodic Table separates the Metals on the left with the non-metals to the right of the line.

The Periods

The horizontal rows are called Periods.

Period 3 contains Sodium(Na), Magnesium(Mg), Aluminium(Al), Silicon(Si), Phosphorus(P), Sulphur(S), Chlorine(Cl) and Argon(Ar).

The Periods represent the energy shell these atoms outer electrons are located within.

Period 3 elements all have their outer electrons in the third energy level/shell.

Period 2 elements have their outer electrons in the second energy shell, and so on.

The transition metals

These atoms have more complicated electronic configurations. They are found in the long block in the middle of the periodic table:

Properties:

Often used in industry as catalysts.

Useful in making alloys.

They form coloured compounds.

The noble gases and halogens

1. The noble gases

This group contains helium, neon, krypton, xenon and radon.

They are different to elements belonging to other groups due to their resistance to form compounds. The reason behind their unreactivity is their full outer shells that give stability to the atoms.

Properties:

Non-metals.

Gases.

Colourless.

Although they have similar properties they are not identical. For example, as you descend Group 0 the density of the gas increases as does the mass of a single atom.

Group VII - The halogens

Fluorine, chlorine, bromine and iodine make up the family of halogens. The last three being the most common. All the halogens exist as molecules, bonding covalently to their own atoms - adding stability as they complete their full outer shells.

Properties:

They all form coloured vapours:

Chlorine - green.

Bromine - red/brown.

Iodine - purple.

Melting points and boiling points are relatively low due to molecules been held together by weak inter-molecular forces. As you descend the group the melting and boiling point increases as the attraction between molecules gets larger.

As with the noble gases, the halogens do have similar properties but not exactly the same ones. For example, the reactivity of the element decreases as you descend the group.

This can be seen if we observe the reaction between iron wool and the different halogens.

So why are the halogens reactive?

The answer lies in the electronic configurations and specifically their outer shell electron configurations. The halogens need only gain 1 electron from another atom to gain more stability.

Fluorine is the most reactive since the electron it is attempting to attract is coming into a shell closest to the positive nucleus. Greater attraction means that it is easier to gain an extra electron - therefore it is the most reactive.