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Non-metals Flashcards

1
Q

How are non-metals bonded?

A

Covalently

note: four covalent bonds form when pairs of electrons are localised in CH4 between the hydrogen and carbon nuclei.

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2
Q

Are there chemical bonds between the molecules themselves?

A

NO. There are only chemical bonds between the atoms in each molecule.

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3
Q

What are the two types of molecules non-metal compounds form?

A

Simple molecular and giant molecular (giant covalent molecular)

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4
Q

Properties of simple molecular substances?

A

Can be all states that have relatively low melting points and boiling points an do not conduct electricity as:
The attractive forces between the molecules are weak
The molecules do not carry an overall charge.

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5
Q

What are allotropes

A

Different forms of the same element in the same state. Different arrangements of atoms.

Each allotrope is regarded as a different substance.

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6
Q

How come C has three allotropes?

A

Because C has the ability to bond with up to four other C atoms.

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7
Q

Properties of non-metals

A

Not good thermal or electricity (except graphite) conductors.
Not lustrous
Not malleable
Not ductile.

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8
Q

Properties of Diamond

A
  • Very hard due to 3D tetrahedral arrangement of C atoms covalently bonded. Hardest natural substance known.
  • Colourless, transparent and crystalline.
  • solid at room temp. Very high melting point due to many strong covalent bonds need to be broken.
  • Insoluble in all liquids
  • Does not conduct electricity due to the lack of delocalised electrons so it’s a perfect insulator.

Reaction with oxygen?
C + O2 -> CO2 when O2 is plentiful
2C + O2 -> 2CO when limited supply of oxygen

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9
Q

What is diamond used for?

A
  • Abrasive (a substance used for grinding, polishing, or cleaning a hard surface) for cutting instruments
  • Jewellery as the hardness of diamond allows faces to be cut into the diamond’s surface
  • as micro-bearings, where tiny diamonds reduce friction between surfaces.
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10
Q

Examples of C in combined form

A

CaCO3 as limestone and marble
CO2 in atmosphere
Organic molecules eg methane and ethanol
Coal

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11
Q

How are the atoms bonded in diamond?

A

Each C atom bonded covalently to four other C atoms. This arrangement forms a continuous, 3D network of C atoms throughout the solid.

This very strong, rigid 3D network of C atoms is what makes it the hardest naturally occurring substance.

Arranged into giant, regular lattice

This forms a giant covalent structure

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12
Q

Draw diagrams of bonding in allotropes of C

A

In notebook

Scipad 69

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13
Q

Properties of graphite

A
  • Soft, slippery feels as sheets of graphite can move over each other due to the weak forces holding the sheets together. The forces are easily broken. Layers of hexagonal rings slide over each other.
  • Insoluble in all liquids.
  • Can conduct electricity because the C atoms only bond to 3 other C atoms. A delocalised e- from each C atom is able to carry a current by moving through the graphite as a charged particle. (the only non-metal that conducts electricity)
  • high melting point. Solid at room temp as the covalent bonds in the layers of atoms in the graphite structure requires a large amount of energy for the network to be broken.
  • most stable form of C
  • grey-black, shiny
  • relatively unreactive

Reaction with oxygen:
C + O2 -> CO2 when O2 is plentiful
2C + O2 -> 2CO when limited supply of oxygen

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14
Q

What is graphite used for?

A
  • Pencils (mixed with clay)
  • Dry lubricant
  • Used in batteries as electrodes due to good electrical conductivity and low reactivity.
  • Can be incorporated into fishing rods, bike frames and golf clubs due to strong covalent bonds between the atoms.
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15
Q

Bonding between atoms in graphite

A

The bonding between C atoms is strong because they are covalently bonded to each other. 3 e- per C atom is used for covalent bonding. The fourth e- bonds very weakly with other C atoms present in layers above and below the C atom to which the e- belongs.

In each layer, rings of 6 C join together to make flat sheets.

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16
Q

Properties of buckminsterfullerene (buckyball)

A
  • C60 molecule
  • very stable
  • able to withstand high temperatures and pressures
  • Low melting point
  • soft and slippery are buckyballs are DISCRETE MOLECULES that are able to move past neighbouring molecules.
  • soluble in organic solvents.
  • low chemical reactivity.
  • yellow powder which turns pink when dissolved in organic solvents such as toluene.
  • solid at room temp: enough intermolecular forces to ensure it is solid at room temp
  • NOT a conductor because no way for these electrons to move between molecules
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17
Q

Bonding in buckyballs

A

C atoms are single bonded to three others, leaving one delocalised electron. Some atoms arrange themselves into 20 hexagons. Other C atoms for single bonds with 3 other C atoms and form 12 pentagons. There are delocalised e- that are able to carry electrical current.

Fullerenes consist of pentagons, hexagons allows it to be bent into spheres, ellipses or cylinders. However, with buckyballs a spheroid is formed. Buckminsterfullerene is the simplest and best-known buckyball.

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18
Q

Uses of buckyballs

A

Their hollow structure allows for other atoms to be housed within them (called doping). Doped with certain metal atoms = superconductor properties
By adding other atoms they can be made into electrical insulators, conductors, semiconductors or even superconductors.

  • using them as super conductors
  • lubricants due to soft, slippery nature.
  • catalysts
  • in fuel cells
  • drug delivery systems where an active drug is encapsulated by the stable C60 molecule and is delivered intact to the target cells in the diseased part of the body as small size (microscopic) and low chemical reactivity.
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19
Q

Carbon nanotubes

A

Structurally similar to buckyballs, but each atom is bonded in a curved sheet that forms a hollow cylinder.

  • high strength
  • high electrical conductivity: because delocalised e- can move along the tube.
  • high heat conductivity
  • high ductility
  • relatively low chemical reactivity
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20
Q

Describe graphene

A

A thin layer of pure C. It is a single, tightly packed layer of C atoms that are bonded together in a hexagonal honeycomb lattice.

Layers of graphene stacked on top of each other form graphite (in simple terms but not quite). Properties slightly different

  • Thinnest compound known to man at one atom thick.
  • Lightest material known
  • Strongest compound discovered
  • Best conductor of heat at room temp and best conductor of electricity known

Uses:

  • ability to replace indium based electrodes in organic LED which reduce the cost of display screens in mobile devices and make them recyclable.
  • Li-ion batteries that contain graphene will recharge faster
  • Uses in objects needing strong strength:weight ratios eg airplane parts
  • Solar cells and fuel cells which will decrease the cost of these technologies.
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21
Q

What molecules does oxygen form?

A

Simple covalent diatomic molecules. O2. each O atom is double bonded to the other.

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22
Q

Properties of Oxygen

A
  • fairly reactive and combines readily with many metals and non-metals, with the release of heat and light energy (ie they burn) in a simple combination/combustion reaction.
  • Metal oxides are ionic. Non-metal oxides are covalent.
  • gas at room temp
  • colourless
  • odourless
  • insoluble (slightly soluble)
  • denser than air
  • 2,6 electron arrangement
  • 21% of the atmopshere
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23
Q

Uses for O2

A
  • Used in the smelting of iron. Injected into molten iron ore to remove sulphur and impurities
  • welding tools: increases flame temp. Use pure O2 instead of air
  • respiration in hospitals.
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24
Q

Describe ozone

A

An allotrope of oxygen.

  • triatomic molecule.
  • Central atom double bonded to one atom and single bonded to the other.
  • less stable than O2
  • formed from O2 by the action of UV light and atmospheric electrical discharges.
  • Human activities have doubled the level of ozone in the troposphere over the last century
  • The highest levels of ozone are found in the stratosphere, called the ozone layer.
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25
Q

Properties of ozone

A
  • Pale blue gas with a strong, distinctive smell
  • ozone in the troposphere is a pollutant, one of the main components of smog.
  • High levels of ozone can be harmful to people, animals and plants.
  • Causes eye irritation and problems with the respiratory system, including asthma, bronchitis and heart disease.
  • Damages plant tissues in crops and forests.
  • can break down materials such as rubber and nylon.
  • oxidises odours
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26
Q

Balanced symbol equation to show formation of zone from diatomic oxygen

A

3O2 -> 2O3

When energy, in the form of an electric spark is passed through oxygen gas, some molecules of oxygen are converted into molecules of ozone.

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27
Q

How does ozone layer work

A

Sunlight is composed of visible light, UV (UV-A, UV-B and UV-C) and infrared light. Light from the UV part of the spectrum causes cellular damage and is linked to the formation of cancers. Increased UV levels also damage plant cells and processes.

The ozone layer absorbs UV radiation from the sun, providing a barrier than prevents most of this radiation from passing to the Earth’s surface (particularly UV-A and UV-B)

The protective role of the ozone layer in the upper atmosphere is so vital that scientists believe life on Earth probably would not have evolved, and could not exist today, without it.

28
Q

Describe the deformation of the ozone layer

A

The ozone hole (over Antarctica) is caused by the catalytic destruction of the ozone molecules by human-made compounds containing halogen atoms, esp compounds called chloroflurocarbons (CFCs). CFCs were widely used as refrigerants, propellants in aerosols and solvents. Their use has been phased out by an international treaty called the Montreal Protocol.

  1. UV breaks a CFCl3 into a CFCl2 molecule and a Cl atom
    CFCl3 + UV ligt -> CFCl2 + Cl
  2. The chlorine atoms reacts with an ozone molecule, taking an oxygen atom with it to form ClO and an oxygen molecule
    Cl + O3 -> ClO + O2
  3. The ClO reacts with a second molecule of O3 to yield another Cl atom and two molecules of O
    ClO + O3 -> Cl + 2O2

This reaction destroys the ozone molecule and recreates the original Cl atom, which can repeat reaction 2 and continue to destroy ozone.
Reaction 2 and 3 creates a repeating cycle that keeps destroying ozone.

29
Q

Why is the break down of ozone by CFC a catalytic destruction

A

A free Cl atom is removed from the CFC. The Cl then reacts with ozone to form ClO, but it is regenerated through the reaction of ClO with O3. The net result of reactions 2 and 3 means that Cl is not depleted.

30
Q

What happens if the ozone hole continues to grow

A

More UV-B will be able to reach the surface of the planet resulting in more skin cancers, the life process of plants will be affected and microscopic organisms such as plankton may not be able to survive.

31
Q

Describe how ozone is used to kill microorganisms

A

-Can be used to remove odours (it oxidises the odours) and kill microorganisms.
Ozone is easily produced by passing oxygen (from air) through an electrical discharge. Oxidation occurs when the ozone molecule splits into an oxygen atom and a diatomic oxygen molecule. It is the oxygen atom that is responsible for its deodorising and disinfectant properties.

As ozone molecules make contact with the cell wall, an oxidation reaction called an “oxidative burst” occurs that creates a hole in the bacteria’s cell wall. The bacterium begins to lose its shape and its cell contents escape. Thousands of ozone collisions occur over only a few seconds resulting in cellular disintegration. The ozone rapidly decomposes to oxygen, leaving no traces of toxic compounds. Cellular death. It is the strong oxidising effect of ozone that is responsible for ozone’s disinfecting properties.

Ozone works in much the same way on viruses, oxidising the protein coat that surrounds them, making them inactive.

Ozone is composed of three Oxygen atoms. One of the atoms is connected to the others weakly and will transfer itself to other substances such as viruses and bacteria, causing them to oxidize by binding itself onto them.

-Used at very low concentrations to purify and disinfect air and water.
Ozone is particularly useful for water treatment as it does not react with water so there is no effect on pH or taste. It also removes pre-existing odours.

32
Q

When is ozone toxic?

A

Ozone is toxic to humans at concentrations above 100 parts per billion. At this level ozone causes damage to mucus and respiratory issues. It can also damage tissues in plants.

Any commercial applications must use ozone at concentrations lower than this in order to be safe to handle.

33
Q

Disadvantages of using ozone

A

Cannot be stored and transported like other industrial gases because it quickly decays into diatomic oxygen, meaning ozone must be produced on site.

  • toxic at high concentrations
  • must be handled carefully
34
Q

Advantages of using ozone

A
  • can be easily produced and cost effective
  • breaks down to form harmless by-products.
  • kills microbes much faster than other oxidising agents. Although Cl kills microbes, ozone leaves no harmful or smelly-by products as the only residue is oxygen gas
35
Q

3 commercial uses of ozone

A
  • sterilising hospital equipment
  • antimicrobial treatment of foods
  • water treatment.
36
Q

What is oxidation?

A

A substance is oxidised, ie combines with oxygen and/or loses electrons

37
Q

Properties of Chlorine

A
  • forms simple covalent, diatomic molecules. Cl2. Each atom singly bonded to the other
  • Reactive and combines readily with most metals to produce chlorides. eg 2Na + Cl2 -> 2NaCl
  • very toxic
  • gas at room temp
  • light yellowy green
  • pungent odour
  • soluble in water. HCl and HOCl (hypochlorous acid) are produced. Cl2 + H2O -> HCl + HOCl
  • denser than air
  • 2,8,7
  • turns damp blue litmus paper red because its acidic, and then bleaches.
  • when a mixture of chlorine and hydrogen gas is heated hydrogen chloride gas is produced. H2 + Cl2 -> 2HCl
38
Q

How is chlorine used in bleach and water purification?

A

NaClO based bleaches are found in many household cleaners. Adding NaClO to water liberates chlorine. Chlorine has bleaching (whitening) properties.

Chlorine changes the pH of water and quickly reacts with microbial cells to irreversibly denature and destroy many pathogens.

39
Q

Describe chlorine and water treatment (kill bacteria in drinking water and swimming pools)

A

Water is disinfected to kill harmful micro-organisms and to make it safe to drink by using either gaseous Cl2 or sodium hypochlorite (NaClO). In both cases the active agent is hypochlorous acid (HOCl). It is thought that chlorine acts in a similar manner to that of ozone, ie the chlorine destroys the cell wall causing the micro-organism to release vital cell contents. Once inside the micro-organism, chlorine also denatures many of the cellular proteins meaning it cannot grow, replicate or repair itself.

Flocculation
Chemicals are added that cause sediments to clump together. These larger particles fall to the bottom of the sedimentation tank and are removed.

Filtration
Filters are used to clarify the water further by removing any remaining small particles. The filters are cleaned regularly.

Disinfection
Chlorine gas or sodium hypochlorite are used to kill microbes.
Cl2 + H2O -> HCl + HOCl
The hydrochloric acid releases H+ ions, making the solution acidic:
HCl -> H+ + Cl-
It is the oxidising effect of the hypochlorous acid, HOCl, that kills bacteria.

Some residual chlorine remains in the water to prevent reinfection in the pipes between the plant and house taps. The pH is raised to 8 to prevent corrosion of the water pipes by adding a neutralising alkali.

40
Q

Two other uses of chlorine disinfection

A
  • Food treatment (ie washing salad leaves)
  • Disinfecting equipment
  • As a bleach in the pulp and paper industry - chlorine is used to bleach wood pulp before the pulp is used to make paper.
  • in making the monomer of the plastic PVC (polyvinylchloride)
41
Q

What is responsible for bleaching and sterilising action?

A

When chlorine is added to NaOH solution, sodium hypochlorite is formed
2NaOH + Cl2 -> NaClO + NaCl + H2O

NaClO is commonly found in chlorine bleaches. It is the hypochlorite ion, OCl-, that is responsible for sterilising and bleaching action.

42
Q

Properties of nitrogen

A
  • Colourless
  • Odourless
  • Tasteless
  • Diatomic gas
  • makes up 78% of the Earth’s atmosphere
  • present in all living tissues. Essential element for life because it occurs in DNA
  • forms N(2-) ion
  • exists as N2 gas. Two N atoms are bonded by a triple bond.
  • slightly less than air
  • largely insoluble in water.
  • does not ignite/burn nor allows other substances to burn in it.
  • fairly unreactive due to strong triple covalent bond between the two N atoms
  • lightning can cause a small amount of atmospheric nitrogen and oxygen to react to produce nitric oxide N2 + O2 -> 2NO. NO combines with more oxygen to produce nitrogen dioxide. 2NO + O2 -> 2NO2
  • NO2 are also produced by the burning of fossil fuels in motor vehicles and industry.
  • in the presence of sunlight, nitrogen oxides combine with other atmospheric pollutants to produce photochemical smog - a toxic mixture of chemicals and particulates, which forms a brown haze over cities and industrial areas. This smog can irritate the breathing passages, and be harmful to people with heart and lung problems.
43
Q

How is nitrogen prepared?

A

Fractional distillation of air. Boiling point of -196°C

44
Q

What is nitrogen used for?

A
  • used as an inert atmosphere for some processes such as refrigerant and to prevent chemical reactions.
  • Food processing/packaging to keep foods eg chips fresh and to prevent crushing
  • coolant or remove warts (liquid nitrogen) because of its low temp (bp -196°C)
  • used in large quantities to produce ammonia gas by the Haber Process. This is then used to make nitric acid, fertiliser, dyes and explosives.
45
Q

Describe the nitrogen cycle

A

Living organisms need nitrogen to make proteins that are needed for growth and repair of new tissue. Nitrogen gas (N2) is most abundant gas in the atmosphere. However N2 cannot be used by plants and animals. Plants need nitrogen in the form of nitrates (NO-3 ions). Animals get the nitrogen they need by eating plants and other animals.

Nitrogen-fixing bacteria convert nitrogen gas from the air into nitrogen containing compounds such as nitrates. Nitrogen-fixing bacteria are found in the soil and also in the root nodules in the roots of legumes. Legumes are a family of plants that includes peas, beans and clover.

Fungal and bacterial decomposers convert nitrogen-containing compounds (eg amino acids) found in dead material and animal excrement into nitrates.

Denitrifying bacteria convert nitrates into nitrogen gas that is released back into the atmosphere.

Nitrogen containing compounds are also found in fertilisers and are used to promote plant growth.

46
Q

How is ammonia made?

A

Made from nitrogen and oxygen

Can be prepared in the lab by heating an ammonium compound (eg ammonium chloride) with an alkali such as calcium hydroxide. The gas produced can be dried using a desiccant such as calcium oxide as it separates water from the ammonia. CaO + H2O -> Ca(OH)2
-boiling tubes must be dry as ammonia dissolves in water to form ammonia solution NH4OH

47
Q

Properties of ammonia

A
  • very soluble in water and forms ammonia solution (NH4OH). The NH4OH then dissociates into NH4+ and OH-, turning water basic.
  • most important use of ammonia in NZ is the production of fertilisers for agricultural purposes.
  • pungent smelling
  • colourless
  • less dense than air
  • turns damp red litmus paper blue.
  • high solubility in water shown by fountain experiment
  • base. Neutralised by acids to produce ammonium salts. eg NH3 + Cl -> NH4Cl
48
Q

Draw lewis diagram of ammonia

A

OK

49
Q

Why can ammonia be collected in an upside down test tube?

A

Because it is less dense than air. Collected by downward displacement of air where ammonia goes up and air comes down.

50
Q

Balanced symbol equation for ammonium chloride and calcium hydroxide

A

2NH4Cl + Ca(OH)2 -> 2NH3 + 2H2O + CaCl2

51
Q

Explain the Fountain experiment

A

A small amount of water is sometimes needed to start the experiment. This dissolves some of the gas and creates a vacuum. This sucks water up the tube and into the test tube. More ammonia dissolves, causing a greater vacuum so more water is suck up and so on.

Creates a pressure difference as it is soluble, causes the water to be pushed up.

52
Q

Outline the Haber Process

A

Natural gas is the source of hydrogen, and air is the nitrogen source. A catalyst is used in some stages to speed up the rate of reaction - instead of using higher temperatures. The process involves several stages:

  1. Preparation of a gas mixture of hydrogen and nitrogen gases (synthesis gas) in the correct proportions.
    Hydrogen is produced from steam and natural gas. Steam reforming
    CH4 + H2O -> CO + 3H2 750°C using a nickel catalyst
    Air is added (nitrogen and oxygen):
    4N2 + O2 + 2H2 -> 2H2O + 4N2
    1100°C, using a nickel catalyst
  2. Removal of CO and CO2
  3. formation of ammonia by the Haber Process
    N2 + 3H2 -> 2NH3 + heat
    400°C and high pressure, iron catalyst.
53
Q

What is ammonia used for?

A

Making nitric acid (to manufacture explosives)
Agricultural chemical to replace nitrogen in the soil removed by intensive plant growth
Manufacture of synthetic fibres, such as nylon
Refrigerant

54
Q

Properties of sulphur

A
  • solid at room temp
  • pale yellow crystalline
  • low melting point - weak intermolecular force
  • insoluble in water
  • acidic oxides
  • S8, ring arrangement. Rhombic crystalline form. When heated to 96-119°C it changes to monoclinic.
  • mainly used for the production of sulfuric acid as it has many industrial uses as well as being a lab chemical.
55
Q

Reaction of sulphur and oxygen and properties of SO2

A

burns with blue flame to produce SO2
S + O2 -> SO2

  • strong odour and irritates the throat and lungs
  • produced by industrial processes and can dissolve in rainwater to produce acid rain which is corrosive to metals and carbonates
  • used as a preservative because of its anti-oxidant effect on micro-organisms
  • bleaching agent in the wood pulp and paper industry.
  • colourless gas
  • denser than air
  • dissolves readily in water
  • will not burn in oxygen, and does not support combustion.
  • acidic
56
Q

What are the natural sources of sulphur

A

Sulphur beds (under ground deposits)
Volcanoes
Metal sulfides
Sulphur impurities in natural gas and petroleum

57
Q

Sulphur dioxide and the environment

A

Fossil fuels eg coal, oil and natural gas contain varying amounts of sulphur. When these fuels are burned, sulphur dioxide is released into the atmosphere. The SO2 molecules dissolve in water in the clouds, forming sulphurous acid (H2SO3). This lowers the pH of the rainwater making it more acidic than normal.

The acidic rainwater corrodes buildings and lowers the pH of the soil, thus killing plants and soil organisms, and acidifies the water in lakes and rivers, harming aquatic creatures.

Some methods that reduce SO2 release into the atmosphere:

  • remove the sulphur and its components before burning
  • dissolving SO2 gas out before releasing waste products into the air
  • burning restrictions
  • higher chimney stacks so waste is released higher up into the atmosphere. This dilutes it.
58
Q

What happens when sulphur is heated?

A

Rhombic sulphur melts to form monoclinic sulfur. Yellow liquid of low viscosity. As this liquid is heated, it darkens and the viscosity increases. Thickens til 195°C when it can’t be tipped from the test tube. More heating makes it turn almost black and the viscosity increases again. Sulphur boils at 444°C to form a red-brown vapour.

To make plastic sulphur, pour boiling liquid into cold water. Rapid cooling doesn’t allow the chains to reform into rings so polymerises into a rubbery substance.

59
Q

Uses of sulphur

A
  • makes sulfuric acid by the contact process
  • if some is added to rubber, it makes it harder and more elastic
  • used for manufacture matches, fireworks and in medicines.
  • lime/sulphur is mixed in a solution of calcium hydroxide and it is sprayed on leaves of plants to act as fungicide.
60
Q

Sulphuric vs suphurous

A

Sulphuric is stronger ie pH lower.

Sulphuric forms salts called sulphates whereas sulphurous forms salts as sulphite

61
Q

Describe sulphur and food preservation

A

Food is preserved by either inactivating micro-organisms or by inhibiting their growth rate. SO2 is a REDUCTANT ie it removes oxygen. When oxygen is removed from a plant cell wall, the wall softens and helps food to dry more easily. Micro-organisms need water in order to survive. Therefore, by drying food out it can be preserved and stored for longer periods.

SO2 removes oxygen from micro-organisms, creating an environment in which they cannot carry out their life functions. The lack of microbial activity means food is less likely to spoil.

SO2 is acidic in solution, causing the pH to decrease creating an environment in which micro-organisms cannot live. The enzymes within micro-organisms can only carry out their function within a very narrow pH range. Altering pH causes the enzymes to denature, Denaturing enzymes means micro-organisms cannot carry out life processes that lead to the spoiling or decolourising of food.

SO2 cannot be added directly to foods. Compounds that release SO2 are added. Sulphites are compounds that contain the ion SO3(2-). When some sulphite compounds are added to water or acids they release SO2

62
Q

Which compound when added to water or acid release SO2?

A

Sulphites. Contains SO3(2-)

63
Q

Describe the Contact Process

A
  1. SO2 is produced from sulphur or a mineral containing sulphur. Sulphur is burned in air. S + O2
  2. SO2 is converted to SO3. SO2 is mixed with air (oxygen) and passed over a catalyst of vanadium pentoxide. 2SO2 + O2 -> 2SO3. Temp of 400°C to melt the catalyst to increase its efficiency.
  3. SO3 is dissolved in H2SO4 to make oleum, H2S2O7. SO3 is absorbed in pure sulfuric acid. SO3 + H2SO4 -> H2S2O7. Carried out at room temp. Directly dissolving SO3 in water is not practical due to the large amount of heat given off in the reaction between SO3 and H2O. The reaction forms corrosive droplets, which would fill the surroundings.
  4. The oleum is diluted with water to produce pure H2SO4. The oleum is carefully mixed with water in the correct proportions to produce sulfuric acid. H2S2O7 + H2O -> 2H2SO4. The process gives off heat so the mixture must be cooled

Manufactured in industrial plants made of steel components. Pure sulfuric acid does not attack metals, so no corrosion of the plant components by the sulphuric acid.

64
Q

Properties of sulfuric acid

A

-acidic properties when in aqueous solution
eg H2SO4 -> 2H+ + SO4(2-)
-strong acid and ionises completely in water
-turns blue litmus red
-releases hydrogen gas when Mg (and other reactive metals) are added: H2SO4 + Mg -> MgSO4 + H2
-neutralises to form salts. eg H2SO4 + NaOH -> Na2SO4 + H2O
-dehydrates other chemicals
-oxidises (removes electrons) metals and non-metals
-sulfonates organic molecules (to make detergents)
-acts as catalyst in many chemical reactions

65
Q

Uses of sulfuric acid

A

-production of fertilisers eg superphosphate and ammonium sulfate. Superphosphate is very important fertiliser, used to supply P to soils. It is calcium dihydrogen phosphate mixed with calcium sulfate. It has a much greater solubility in water than rock phosphate. Made by treating rock phosphate with concentrated sulfuric acid

Ca3(PO4)2 + 2H2SO4 -> Ca(H2PO4)2 + 2CaSO4

  • used in manufacture of chemicals eg in making HCl, HNO3, dyes and pigments, explosives and drugs
  • petroleum refining to wash impurities out of gasoline and other refinery products
  • processing metals eg cleaning iron and steel before plating with tin or zinc
  • manufacture of synthetic fibres eg terylene and rayon
  • used as the electrolyte in lead-acid storage battery commonly used in motor vehicles (the acid for this use, containing 1/3 H2SO4, is often called battery acid)
66
Q

Explain H2SO4 in batteries

A

Concentrated sulfuric acid can be stored in metal containers because it is unreactive in its molecular form. However, when even a small amount of water is added it will react vigorously with metals and other chemicals. When sulfuric acid reacts with water it dissociates and forms ions. The formation of the H+ ions makes the solution very reactive.

Dissociation also changes the electrical conductivity of the solutions. When sulfuric acid dissociates it releases ions that are able to carry an electrical current. Substances that dissociate into ions when placed in water are called “electrolytes”. If they dissociate completely they are good electrolytes.

When the battery is discharging, the Pb at the Pb electrode loses two e- and becomes Pb2+. The Pb2+ then reacts and bonds ionically with the negative sulfate ions in solution to form PbSO4. The PbO2 electrode then gains the two e-, becoming PbO2 + 2e-. This negative charge on the electrode attracts the positive H+ ions, meaning that PbO2 + 2e- + 4H+ -> Pb + 2H2O, so lead is formed. The concentration of the acid decreases when it’s discharging

When the battery is charging, Pb turns back into PbO2 as it loses two electrons and reattracts the O2. This means that H+ ions are released into the solution. The PbSO4 then gains two electrons, meaning that the Pb2+ would become lead again, and that the SO4(2-) would be released back into the solution. The SO4(2-) then rebonds with the H+ ions to form H2SO4. The concentration of the acid increases when it’s charging.

67
Q

Examples of carbon compounds with localised and delocalised electrons

A

In diamond, all 4 valence electrons of C are localised between the atoms in covalent bonding. Because the movement of electrons is restricted, an electric current cannot be conducted.

In graphite, each C atom uses 3 of out 4 valence electrons to covalently bond with 3 other C in a plane. Each C atom contributes one electron to the delocalised system of electrons that is also part of the chemical bonding. The delocalised electrons are free to move throughout the plane, meaning that graphite can conduct electricity along the planes of C atoms, but not in a direction perpendicular to the plane

Selected Elements (3 decks)

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