ATOMIC AND PHYSICAL PROPERTIES OF THE PERIOD …chemistrysuzhou.wikispaces.com/file/view/Unit+3... · Web viewYou will see that both melting points and boiling points rise as you

AS ChemistryUnit 3: Inorganic Chemistry

Section 1: Period 3 Part 1: Atomic and Physical Properties of the Elements

Atomic Properties

Task 1

Can you write the electronic configurations of the Period 3 elements in the table below?

Na [Ne] Mg [Ne] Al [Ne] Si [Ne]

P [Ne] S [Ne] Cl [Ne] Ar [Ne]

In each case, [Ne] represents the complete electronic structure of a neon atom.

Atomic radius

The graph shows how the atomic radius changes as you go across Period 3:

Na Mg Al Si P S Cl Ar0

0.05

0.1

0.15

0.2

0.25

Element

Ato

mic

radi

us (n

m)

Task 2

An atomic radius is a measure of the distance from the nucleus to the bonding pair of electrons.

i) What is the general trend in atomic radius across Period 3?

Cambridge A-level Centre


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ii) Can you use your knowledge of atomic structure to explain this trend?......

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iii) Why do you think the value for argon does not follow the expected trend?....................................................................................................................................

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Ionic radius

The radii of the ions also decrease across a Period, but it must be remembered that elements on the left form cations and elements on the right form anions.

Cations Anionsion Na+ Mg2+ Al3+ Si4+ P3- S2- Cl-ionic radius

0.098 0.065 0.045 0.038 0.212 0.190 0.181



(nm)

Task 3

The radii of the cations are all smaller than that of the corresponding atoms, but the radii of the anions are all larger. Can you explain why?

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First ionisation energy

This has already been covered in Unit 1. You should be able to use your scientific knowledge to explain in detail the trend shown in the graph below:


200400600800

1000120014001600

Element

Firs

t ion

isatio

n en

ergy

(kJ m

ol-1

)

Remember, first ionisation energy is governed by:



the charge on the nucleus; the distance of the outer electron from the nucleus; the amount of screening by inner electrons; whether the electron is alone in an orbital or one of a pair.

Electronegativity

Task 4

Can you write a definition for the term ‘electronegativity?

Electronegativity is .................................................................................................................

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The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7.

The trend

The Electronegativity value for chlorine is given. Can you sketch the rest of the graph?


0.5

1

1.5

2

2.5

3

3.5

Element

Paul

ing

elec

tron

egati

vity

val

ue

Explaining the trend



The trend is explained in exactly the same way as the trend in atomic radii.

As you go across the period, the bonding electrons are always in the same level – level 3. They are always being screened by the same inner electrons.

All that differs is the number of protons in the nucleus. As you go from sodium to chlorine, the number of protons steadily increases and so attracts the bonding pair more closely.

Task 5

Why is no value included for argon?....................................................................................



Physical Properties

Here we look at the electrical conductivity, and the melting and boiling points of the elements. To understand these, you first have to understand the structure of each of the elements.

Task 6

Use your knowledge from Unit 1 to complete the table below.

Na Mg Al Si P S Cl Ar

Structure

Type of element

Bonding

Formula

Type of force broken on melting/boilingDoes the element conduct electricity?



Task 7

Use the information in the table above to answer the following questions.

1. (a) Explain why electrical conductivity decreases across Period 3 from sodium to phosphorus.

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(3)

(b) The table below shows the melting temperatures, Tm, of the Period 3 elements.

Element Na Mg Al Si P S Cl Ar

Tm/K 371 923 933 1680 317 392 172 84

Explain the following in terms of structure and bonding.

(i) Magnesium has a higher melting temperature than sodium.

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(2)

(ii) Silicon has a very high melting temperature.

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(2)



(iii) Sulphur has a lower melting temperature than magnesium.

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(2)

(iv) Argon has a lower melting temperature than chlorine.

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(3)(Total 14 marks)

ReferencesA-level Chemistry pages 197-207Chemistry in Context pages 42-47, 169-176

Learning ObjectivesCandidates should be able to:

describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements (see the Data Booklet)

explain qualitatively the variation in atomic radius and ionic radius interpret the variation in melting point and in electrical conductivity

in terms of the presence of simple molecular, giant molecular or metallic bonding in the elements

explain the variation in first ionisation energy.



Section 1: Period 3 Part 2: Reactions with oxygen

Reactions of Period 3 Elements with oxygenAs we move across a period of the Periodic Table from left to right, we can see that there are small, but regular changes in atomic structure. These small changes can produce very big variations in the chemical and physical properties of the elements and their compounds. The pattern is then repeated as you go across the next period. The occurrence of periodic patterns is called Periodicity. The physical properties of the Period 3 elements have been discussed previously. Here we look at their chemical properties. Task 1

Can you complete the table below?Group number 1 2 3 4 5 6

Element in Period 3

Nuclear charge

[Ne] electronic configurationTrend in Atomic radius

Trend in 1st ionisation energyTrend in electronegativityFormula of oxide/s

The reactivity and properties of elements depend upon a combination of things: nuclear charge, size of the atom, the number of outer electrons and the amount of shielding, and these will also help to explain and predict the properties of many compounds. It is therefore very useful if you are familiar with the trends shown in the table above. You should not, however, lose sight of your general chemistry knowledge.

Task 2

Can you complete the notes below by adding the missing words and writing the appropriate balanced equations?



With oxygen

The solid elements in Period 3 all burn in air or oxygen when ignited. Sodium burns with a __________ flame, forming the oxide:

Magnesium, aluminium, silicon and phosphorus burn when ignited, emitting a very bright white light and white smoke of the oxides:

These reactions are all very exothermic.

Sulphur burns with a __________ flame but much less vigorously than the elements above, to form the pungent, colourless gas ____________________:

In an excess of pure oxygen, some SO3 is also formed. This utilises the highest oxidation state of sulphur.

Task 3

Can you complete the table below?

Na2O MgO Al2O3 SiO2 P4O10 SO2Tm/K 1548 3125 2345 1883 573 200BondingStructure

Use your scientific knowledge to explain the changes in melting point.

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Acid-base properties of the oxides of Period 3 Elements

The table below shows the change in pH for solutions of the oxides across Period 3. There is an evident trend for alkaline oxides → acidic oxides as the bonding changes from ionic to covalent, but it is masked by the change in solubilities. A substance will only change the pH of water if it dissolves.

Task 4

Can you write balanced equations for the reactions of each oxide with water to give the appropriate pH value.

Oxide Reaction with water pHNa2O 14

MgO 9

Al2O37

SiO27

P4O100

SO23

SO30

Ionic oxides

The oxide ion is too highly charged to exist on its own in water. It attracts water molecules and hydrolyses to form OH-(aq) ions:

O2− + H2O → 2 OH−

Na2O: very soluble in water → lots of oxide ions present and therefore many OH-(aq) ions will be formed.

MgO: less soluble in water due (in part) to higher lattice energy → less oxide ions and less OH-(aq) ions.

Al2O3: insoluble → no oxide ions will be present and hence no OH-(aq) ions will be formed. Aluminium oxide is an ionic solid with some covalent character. It is amphoteric; it will act as both an acid and a base.



BASE: Al2O3 + 3H2SO4 → Al2(SO4)3 + 3H2O

ACID: Al2O3 + 2NaOH + 3H2O → 2NaAl(OH)4

Covalent oxides

SiO2: insoluble in water → no ions produced.

P4O10, SO3: these acidic oxides rapidly dissolve in water and give the corresponding acids. They are very soluble, so many H+(aq) ions are formed, giving a low pH.

SO2 is less soluble so a weaker acid is formed.

ReferencesA-level Chemistry pages 209-211Chemistry in Context pages 177-184


describe the reactions, if any, of the elements with oxygen to give Na2O, MgO, Al2O3, P4O10, SO2 and SO3.

state and explain the variation in oxidation number of the oxides. describe the reactions of the oxides with water. describe and explain the acid/base behaviour of oxides and

hydroxides, including, where relevant, amphoteric behaviour in reaction with NaOH and acids.



Section 1: Period 3 Part 3: Reactions with chlorine

All of the chlorides of the period 3 elements can be formed by direct combination with chlorine but your syllabus is only interested in the reactions of Na, Mg, Al, Si and P.

Task 1Use the information on pages 47-49 and 179-180 of ‘Chemistry in Context’ and pages 209 and 211-212 of ‘AS level Chemistry’ to complete the table below:

Element Na Mg Al Si PDescription of reaction with chlorine

Formula of chloride/s

Oxidation state of period 3 element

State of chloride at r.t.p.

b.pt. of chloride (oC)

Structure of chloride

Bonding in chloride


AS Chemistry Unit 3: Inorganic Chemistry

Task 2

Read the information sheet ‘Aluminium chloride’ and answer the question below.

i. At room temperature aluminium chloride is considered to be ionic. What two pieces of evidence seem to contradict this?

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ii. At temperatures above ~200oC aluminium chloride exists as a dimer, Al2Cl6. This molecule contains both ordinary and dative covalent bonds. Draw a diagram to show the two types of bonding.

Reaction of the period 3 chlorides with water

As Period 3 is crossed, the reactions of the chlorides with water become increasingly more violent.

Ionic chlorides usually dissolve in water to form neutral solutions containing the hydrated parent ions.

Covalent chlorides are hydrolysed by water to form acidic solutions containing HCl.

Task 3

Can you summarise these reactions in the table below?

Reaction with water pH



7

6/7

3

0

0

ReferencesA-level Chemistry pages 209-210 and 211- 212Chemistry in Context pages 47-49 and 179-180


describe the reactions, if any, of the elements with chlorine to give NaCl, MgCl2, Al2Cl6, SiCl4, and PCl5.

describe and explain the reactions of the chlorides with water.



Section 2: Group 7 Part 1: Properties of the elements

Please note that fluorine will not be included in discussions of Group VII, because its small size and high electronegativity give it some anomalous properties.

Task 1

Use your general scientific knowledge to fill in the gaps in the paragraph below.

The halogens are the __________ in _______7 of the __________ Table. All halogen atoms have 7 __________ in the outer shell. The halogens are the most __________ group of non-metals, and none of them is found __________ in the elemental form. They are all found in __________, often as __________ ions (a singly __________ charged ion, e.g. Br-).

Fluorine and chlorine are the most __________ halogens, bromine occurs in smaller __________, iodine is quite __________ and astatine is an __________ produced, short-lived, __________ element.

All the halogen elements occur as __________ molecules, e.g. F2. The atoms are linked by a single __________ bond.

In compounds, a halogen atom can attain stability by: __________ an electron from a __________ atom to form a halide ion in

an __________ bonded compound,

sharing an electron from another atom in a __________ bonded compound.

Task 2

Draw electron dot-cross diagrams for NaCl and HCl in the space below.

Task 3

Use the information in your textbooks to complete the table below.



Element

Electronic structure

Appearance at r.t.p.

State at r.t.p.

Melting point (oC)

Boiling point (oC)

Colour in water

Colour in hexane

FluorineChlorine BromineIodine

Task 4

From your knowledge of the structure and bonding of the halogens, explain why they are more soluble in organic solvents than in aqueous solution.

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Trends in Melting Point and Boiling Point



You will see that both melting points and boiling points rise as you descend the Group.

Task 5

Use your scientific knowledge to explain the trend in volatility of the halogens........................................................................................................................................................

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Trends in chemical properties

Task 6

Watch the short video extract. For each reaction write down your observations.

Halide ion in aqueous solutionHalogen Chloride, Cl- Bromide, Br- Iodide, I-Chlorine, Cl2

Bromine, Br2

Iodine, I2

The oxidising power of the halogens

Imagine the reaction between one halogen (chlorine, say) and the ions of another one (iodide ions, perhaps). The iodide ions will be in a solution of a salt like sodium or potassium iodide. The sodium or potassium ions will be spectator ions, and are completely irrelevant to the reaction. In the chlorine and iodide ion case, the reaction would be:

Cl2 + 2I- 2Cl- + I2



The iodide ions have lost electrons to form iodine molecules. They have been __________.

The chlorine molecules have gained electrons to form chloride ions. They have been __________.

This is obviously a redox reaction in which chlorine is acting as _______________ agent.

Fluorine

We'll have to exclude fluorine from this descriptive bit, because it is too strong an oxidising agent. Fluorine oxidises water to oxygen and so it is impossible to do simple solution reactions with it.

Chlorine, bromine and iodine

In each case, a halogen higher in the Group can oxidise the ions of one lower down. For example, chlorine can oxidise the bromide ions (in, for example, potassium bromide solution) to bromine:

The bromine appears as an orange solution.

As you have seen above, chlorine can also oxidise iodide ions (in, for example, potassium iodide solution) to iodine:

The iodine appears either as a red/brown solution if you are mean with the amount of chlorine you use, or as a dark grey precipitate if the chlorine is in excess.

Bromine can only oxidise iodide ions to iodine. It isn't a strong enough oxidising agent to convert chloride ions into chlorine. (You have just seen exactly the reverse of that happening.)

A red/brown solution of iodine is formed (see the note above) until the bromine is in excess. Then you get a dark grey precipitate.


Fluorine

Chlorine

Bromine

Iodine

Relative oxidising power


Iodine won't oxidise any of the other halide ions (unless you happened to have some extremely radioactive and amazingly rare astatide ions - astatine is at the bottom of this Group).

To summarise

Oxidation is loss of electrons. Each of the elements (for example, chlorine) could potentially take electrons from something else to make their ions (e.g. Cl-). That means that they are all potentially oxidising agents.

Fluorine is such a powerful oxidising agent that you can't reasonably do solution reactions with it.

Chlorine has the ability to take electrons from both bromide ions and iodide ions. Bromine and iodine can't get those electrons back from the chloride ions formed. That means that chlorine is a more powerful oxidising agent than either bromine or iodine.

Similarly bromine is a more powerful oxidising agent than iodine. Bromine can remove electrons from iodide ions to give iodine - and the iodine can't get them back from the bromide ions formed.

This all means that oxidising ability falls as you go down the Group.

Reactions of the halogens with hydrogen

The halogens all react with hydrogen to give the corresponding hydrogen halide. The relative reactivity of the halogens is well illustrated by these reactions:

Fluorine explodes with hydrogen even in the dark at -200oC. The reaction with chlorine is explosive when exposed to ultraviolet

light. With bromine, the reaction occurs slowly on heating. The reaction with iodine is incomplete on heating.

Task 7

Write equations, including state symbols, for the reactions of chlorine, bromine and iodine with hydrogen gas.



The trend is further illustrated by consideration of the ΔHfθ values:

HCl HBr HIΔHfθ (kJ mol-1) -92 -36 +26

The hydrogen halides formed are all simple molecular compounds which are gaseous at room temperature. As the size of the halogen atom increases down the Group, the H-X bond length also increases. Consequently, the H-X bond enthalpy decreases down the Group and the hydrogen halides become less stable:

hydrogen chloride is stable at 1500oC hydrogen bromide decomposes appreciably at 800oC, and hydrogen iodide decomposes appreciably at 500oC.

If a red hot glass rod is placed in a gas jar containing hydrogen iodide, purple fumes of iodine are seen. This shows that iodide ions are easily oxidised and can act as strong reducing agents.



describe the trends in volatility and colour of chlorine, bromine and iodine.

interpret the volatility of the elements in terms of van der Waals’ forces.

describe the relative reactivity of the elements as oxidising agents. describe and explain the reactions of the elements with hydrogen. describe and explain the relative thermal stabilities of the hydrides. interpret these relative stabilities in terms of bond energies.



Section 2: Group 7 Part 2: Properties of the halides

Here we look at the redox reactions involving halide ions and concentrated sulphuric acid. We use these reactions to discuss the trend in reducing ability of the ions as you go from fluoride to chloride to bromide to iodide.

The Facts

There are two different types of reaction which might go on when concentrated sulphuric acid is added to a solid ionic halide like sodium fluoride, chloride, bromide or iodide. The concentrated sulphuric acid can act both as an acid and as an oxidising agent.

Concentrated sulphuric acid acting as an acid

The concentrated sulphuric acid gives a hydrogen ion to the halide ion to produce a hydrogen halide. Because this is a gas, it immediately escapes from the system. If the hydrogen halide is exposed to moist air, you see it as steamy fumes.

As an example, concentrated sulphuric acid reacts with solid sodium chloride in the cold to produce hydrogen chloride and sodium hydrogensulphate.

All of the halide ions (fluoride, chloride, bromide and iodide) behave similarly.

The reduction of concentrated sulphuric acid

With fluoride or chloride ions

The fluoride and chloride ions aren't strong enough reducing agents to reduce the sulphuric acid.

Whichever way you look at it, all you get is the hydrogen halide!

That isn't true, though, with bromides and iodides.

With bromide ions

The bromide ions are strong enough reducing agents to reduce the concentrated sulphuric acid to sulphur dioxide gas. In the process the bromide ions are oxidised to bromine.



Task 1

Can you write two electron half equations for these steps, then combine them to form the overall redox equation for the reaction.

Task 2

Complete the notes below.

The bromide ions reduce the sulphuric acid to sulphur dioxide gas. This is a decrease of oxidation state of the sulphur from _____ in the sulphuric acid to _____ in the sulphur dioxide.

What you see in this reaction are the steamy fumes of __________ __________ contaminated with the brown colour of __________ vapour. The sulphur dioxide is a colourless gas, so you couldn't observe its presence directly.

With iodide ions

Iodide ions are stronger reducing agents than bromide ions are. They are oxidised to iodine by the concentrated sulphuric acid. The reduction of the sulphuric acid is more complicated than before. The iodide ions are powerful enough reducing agents to reduce it

first to sulphur dioxide then to sulphur itself and all the way to hydrogen sulphide.

Task 3

Can you give the oxidation state of sulphur in each of the species above?

The most important of this mixture of reduction products is probably the hydrogen sulphide.

Task 4


Fluoride

Chloride

Bromide

Iodide

Relative reducingpower


Can you write the half equations for these reactions and combine them to give the overall redox equation?

This time what you see is a trace of steamy fumes of hydrogen iodide, but mainly lots of iodine. The reaction is exothermic and so purple iodine vapour is formed, and probably dark grey solid iodine condensing around the top of the tube.

You won't see the colourless hydrogen sulphide gas, but might pick up its "bad egg" smell if you were foolish enough to smell the intensely poisonous gases evolved!

Summary of the trend in reducing ability

Fluoride and chloride ions won't reduce concentrated sulphuric acid. Bromide ions reduce the sulphuric acid to sulphur dioxide. In the

process, the bromide ions are oxidised to bromine. Iodide ions reduce the sulphuric acid to a mixture of products

including hydrogen sulphide. The iodide ions are oxidised to iodine. Reducing ability of the halide ions increases as you go down the

Group. (Obviously, this is the opposite direction to their oxidising ability!!)

Testing for halide ions

Using silver nitrate solution



Carrying out the test

This test has to be done in solution. If you start from a solid, it must first be dissolved in pure water.

The solution is acidified by adding dilute nitric acid. (Remember: silver nitrate + dilute nitric acid.) The nitric acid reacts with, and removes, other ions that might also give a confusing precipitate with silver nitrate.

Silver nitrate solution is then added.

Task 5

Can you complete the table below to show what you would observe in each case?

ion present observation

F-

Cl-

Br-

I-

All of the precipitates change colour if they are exposed to light - taking on grey or purplish tints.

The absence of a precipitate with fluoride ions doesn't prove anything unless you already know that you must have a halogen present and are simply trying to find out which one. All the absence of a precipitate shows is that you haven't got chloride, bromide or iodide ions present.

The chemistry of the test

The precipitates are the insoluble silver halides - silver chloride, silver bromide or silver iodide.

Silver fluoride is soluble, and so you don't get a precipitate.



Confirming the precipitate using ammonia solution

Carrying out the confirmation

Ammonia solution is added to the precipitates.

Task 6

Can you complete the table below to show what you would observe in each case?

original precipitate observation

AgCl

AgBr

AgI

The chemistry of the test

When the precipitate of silver chloride or silver bromide dissolves, the silver ion forms a complex ion:

E.g. AgCl(s) + 2NH3(aq) → [Ag(NH3)2]+(aq) + Cl-(aq)

The reactions of chlorine with sodium hydroxide

Chlorine is used in the manufacture of bleach, which is widely used as a disinfectant. Bleach is produced by passing chlorine gas up a tower, down which cold, dilute aqueous sodium hydroxide is flowing. The equation for the reaction which takes place in the tower is as follows:

Cl2(g) + 2NaOH(aq) → NaClO (aq) + NaCl(aq) +H2O(l)

Sodium chlorate (I)

Task 7

i. Using oxidation states can you identify which element is oxidised and which is reduced in the equation above?

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ii. What word do we use to describe reactions of this sort?



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With hot, concentrated aqueous sodium hydroxide a similar reaction occurs but this time a mixture of chloride and chlorate (V) ions is formed.

Task 8

Can you write a balanced equation for the reaction above?

The sodium chlorate (V) produced is used as a weedkiller.

Task 9

Use pages 242-243 of your AS level textbook to list the uses of the halogens below. Add any chemical formulae or equations you think may be useful.


Uses of the halogens




describe and explain the reactions of halide ions with o aqueous silver ions followed by aqueous ammoniao concentrated sulphuric acid.

describe and interpret in terms of changes of oxidation number the reaction of chlorine with cold, and with hot, aqueous sodium hydroxide.

explain the use of chlorine in water purification. recognise the industrial importance and environmental

significance of the halogens and their compounds, (e.g. for bleaches; PVC; halogenated hydrocarbons as solvents, refrigerants and in aerosols).



Section 3: Nitrogen and Sulphur

Nitrogen

Use pages 259-264 of ‘AS level Chemistry’ to find the answers to the following questions?

1. Nitrogen gas makes up what percentage of the Earth’s atmosphere?

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2. Draw a dot-cross diagram to show the bonding in a molecule of nitrogen:

3. Nitrogen is used as an inert atmosphere in food storage and during some metalworking processes where oxidation must be prevented. Why do you think nitrogen gas is so unreactive?

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4. Can you list two situations in which nitrogen gas will react with oxygen gas? Why do these reactions occur?

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5. Atmospheric nitrogen is ‘fixed’ as nitrate ions by bacteria in the roots of certain plants such as peas, beans and clover. How can bacteria enable this to happen?

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6. Nitrogen is converted into ammonia in the Haber Process. Can you write a balanced equation for this reaction.

7. Ammonia readily forms ammonium ions through co-ordinate bonds. Can you draw a diagram to represent this below:

Ammonia is acting as what type of species in this reaction?

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8. How would you test for the presence of an ammonium ion?

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9. Much ammonia is converted into salts for use as fertilisers. What is a fertiliser?

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Can you name three of these salts?

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10. Ammonia is also converted into nitric acid. Can you write a balanced equation for the overall reaction and give the conditions needed?

11. There are concerns over the use of nitrate fertilisers as they have been linked to the process of eutrophication. Can you draw a flow diagram for this process below:



Sulphur

Use your scientific knowledge and the information on pages 262-264 of your textbook to help you answer the following questions.

1. The Earth’s atmosphere contains oxides of nitrogen and sulphur dioxide from natural sources. Can you name some of these natural sources?

a. nitrogen oxides........................................................................................................

b. sulphur dioxide........................................................................................................

2. Sulphur dioxide in the atmosphere has been linked to human respiratory problems such as bronchitis and asthma. It is however, used as a food preservative. How does sulphur dioxide preserve food?

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3. Sulphur dioxide is oxidised in the atmosphere by nitrogen dioxide. Write a balanced equation for this reaction.

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4. Acid rain is corrosive and will attack buildings made of limestone and marble. Why are these buildings so vulnerable?

5. What is being done to reduce sulphur dioxide emissions?

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6. Sulphur extracted from crude oil is used to make sulphuric acid in the Contact Process. Can you give four uses for sulphuric acid?

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ReferencesA-level Chemistry pages 259-264


explain the lack of reactivity of nitrogen. describe the

o formation, and structure of, the ammonium iono the displacement of ammonia from its salts.

understand the environmental consequences of the uncontrolled use of nitrate fertilisers.

understand and explain the occurrence, and catalytic removal, of oxides of nitrogen.

explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulphur dioxide.

describe the formation of atmospheric sulphur dioxide from the combustion of sulphur contaminated carbonaceous fuels.

state the role of sulphur dioxide in the formation of acid-rain and describe the main environmental consequences of acid-rain.

understand the industrial importance of sulphuric acid. Describe the use of sulphur dioxide in food preservation.



Section 4: Group 2 Section 1: Atomic and physical properties

Atomic Radius

You can see that the atomic radius increases as you go down the Group.

Explaining the increase in atomic radius

Task 1

Which factors affect the size of the atomic radius?

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Compare beryllium and magnesium:

Be 1s22s2

Mg 1s22s22p63s2

In each case, the two outer electrons feel a core charge of 2+ from the nucleus. The positive charge on the nucleus is shielded by the inner electrons.



This is equally true for all the other atoms in Group 2. Work it out for calcium if you aren't convinced.

The only factor which is going to affect the size of the atom is therefore the number of layers of inner electrons which have to be fitted in around the atom. Obviously, the more layers of electrons you have, the more space they will take up - electrons repel each other. That means that the atoms are bound to get bigger as you go down the Group.

Trends in First Ionisation Energy

First ionisation energy is the energy needed to remove the most loosely held electron from each of one mole of gaseous atoms to make one mole of singly charged gaseous ions - in other words, for 1 mole of this process:

X(g) ® X+(g) + e-

Notice that first ionisation energy falls as you go down the group.

Task 2

Can you explain this trend?



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Trends in Electronegativity

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. All of these elements have a low electronegativity (due to their small core charge).

As the metal atoms get bigger, any bonding pair gets further and further away from the metal nucleus, and so is less strongly attracted towards it. In other words, as you go down the Group, the elements become less electronegative.



As you go down the Group, the bonds formed between these elements and other things such as chlorine become more and more ionic. The bonding pair is increasingly attracted away from the Group 2 element towards the chlorine (or whatever).


Learning ObjectivesCandidates should be able to interpret, and make predictions from, the trends in physical and chemical properties of the elements and their compounds.



Section 4: Group 2 Section 2: Chemical properties

When Group 2 metals react, they are oxidised from a state of 0 to +2, forming M2+ ions. The elements are powerful reducing agents. Group 2 metals become increasingly reactive as you go down the Group. This is because the distance between the nucleus and the outer electrons increases and so they become easier to remove. Beryllium is markedly different from the other members of the group and will not be considered in detail.

Reactions of the metals with water

Beryllium

Beryllium has no reaction with water or steam even at red heat.

Magnesium

Magnesium burns in steam to produce magnesium oxide and hydrogen.

Very clean magnesium has a very slight reaction with cold water. The reaction soon stops because the magnesium hydroxide formed is almost insoluble in water and forms a barrier on the magnesium preventing further reaction.

Task 1

Can you write equations for the reaction of Mg with both cold and hot water?

As a general rule, if a metal reacts with cold water, you get the metal hydroxide. If it reacts with steam, the metal oxide is formed. This is because the metal hydroxides thermally decompose (split up on heating) to give the oxide and water.

Calcium, strontium and barium

These all react with cold water with increasing vigour to give the metal hydroxide and hydrogen. Strontium and barium have reactivities similar to lithium in Group 1 of the Periodic Table.

The equation for the reactions of any of these metals would be:



The hydroxides aren't very soluble, but they get more soluble as you go down the Group. The calcium hydroxide formed shows up mainly as a white precipitate (although some does dissolve). You get less precipitate as you go down the Group because more of the hydroxide dissolves in the water.

Reactions of the metals with oxygen

The group 2 elements all burn in oxygen to form white solid oxides.

2M(s) + O2(g) → 2MO(s)

Beryllium is reluctant to burn unless it is in the form of dust or powder. Its flame colour is white. Beryllium has a very strong (but very thin) layer of beryllium oxide on its surface, and this prevents any new oxygen getting at the underlying beryllium to react with it.

Magnesium burns with a brilliant white flame. The others burn with characteristic flame colours:

Calcium: brick redStrontium: crimson redBarium: apple green

Reactions of the oxides with water

The metal oxides are all basic and each reacts with water, forming a solution of the hydroxide which is alkaline. The reaction of magnesium oxide with water is slow; the other oxides react readily with water.

The solubilities of the hydroxides formed increase down the Group and the solutions become more alkaline (approx. pH9 – 12).

Thermal stability of Group 2 carbonates and nitrates

Thermal decomposition is when a substance breaks down (decomposes) when heated. The more thermally stable a substance is, the more heat it will take to break it down.

You are only expected to describe the trend in thermal stability, not explain it.

The thermal stability increases down the group.

Group 2 carbonates decompose to from the oxide and carbon dioxide. Group 2 nitrates decompose to form the oxide, nitrogen dioxide and oxygen:



Task 2

Can you write equations for the thermal decomposition of calcium carbonate and calcium nitrate?

Uses

Task 3

Compounds of the Group 2 elements have many uses. Use the information spread throughout your textbooks to answer the following questions.

1. Magnesium is the most commonly used metal in Group 2. Can you list 2 uses below?

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................2. Which compounds of magnesium are used for the following? In each

case explain why.

a. Toothpaste and indigestion remedies......................................................................

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b. Furnace linings...............................................................................................................

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3. Which 3 types of rock contain large amounts of calcium carbonate?

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4. Several calcium compounds have common names based on the word ‘lime’. Can you draw a flow diagram in the space below to show the relationship between limestone, lime (or quicklime), slaked lime and limewater. Include the chemical name and formula for each substance.

5. Find as many uses as you can for each of the following compounds of calcium:

a. calcium carbonate.........................................................................................................

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b. calcium oxide.................................................................................................................



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c. solid calcium hydroxide...............................................................................................

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d. calcium hydroxide solution.........................................................................................

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Learning ObjectivesCandidates should be able to

describe the reactions of the elements with oxygen and water. describe the behaviour of the oxides with water. describe the thermal decomposition of the nitrates and carbonates. interpret, and make predictions from, the trends in chemical

properties of the elements and their compounds. explain the use of magnesium oxide as a refractory lining material

and calcium carbonate as a building material. describe the use of lime in agriculture.


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ATOMIC AND PHYSICAL PROPERTIES OF THE PERIOD …chemistrysuzhou.wikispaces.com/file/view/Unit+3... · Web viewYou will see that both melting points and boiling points rise as you