ZON PENA

Stability of Noble Gases

What is Chemical Compound

• A chemical compound is a substance that is formed by more than one elements that bond together chemically in a fixed proportions.
• In periodic table, there are only 118 elements, and about 1/3 of them are synthetic elements.
• Only a few substances exist as element (Not Compound) in nature.
• The table below shows some examples of substance exist as element in nature.

Element exist as monoatomic gas	Element exist as diatomic Molecule(gas)	Element exist as solid
Helium (He) Neon (Ne) Argon (Ar) Krypton (Kr) Xenon (Xe) Radon (Rn)	Oxygen (O2) Nitrogen (N2)	Carbon (Graphite & Diamond) Gold Silver Platinum

• In nature, we can find millions of substances, which means most of the chemical substances exist as compound in nature.
• In short, elements tend to form compound in nature.

Why Elements Tend to Form Compound?

• A compound is formed by 2 or more elements hold together by a force called chemical bond.
• Before studying why elements like to bond together, we need to know why certain elements such as Helium and Neon do not form any bonds with other elements.

Why Noble Gases Don't Form Compound

• In previous chapter, we have discussed that Group 18 elements (Noble Gases) exist as monoatom in nature.
• They are inert in nature and do not react with any other elements (or themselves) to form any chemical compounds.
• In other words, they are chemically very very stable (or chemically very very non-reactive).

Duplet and Octet Electron Arrangement

• The charge on the nucleus and the number of electrons in the valence shell determine the chemical properties of an atom.
• The stability of noble gas is due to their electrons arrangement.
• The diagram above shows the first four elements of Noble Gas.
• We can see that the outer most shell (valence shell) of Helium has 2 electrons. We call this duplet electron arrangement. We should take notes that the maximum number of electrons can be filled in the first shell is 2 electrons, which means 2 electrons in the first shell is considered FULL.
• The valence shell all other Group 18 elements (including Xenon and Radon which is not shown in the diagram) has 8 electrons, and we call this octet electron arrangement.
• When the electron arrangement of an atom is duplet or octet, the energy of the electrons is very low, and it is very difficult (even though it is not impossible) to add or remove electrons from the atom.
• This explain why noble gases are reluctant to react with all other elements.

The Octet Rule

• So far we have learnt that the electon arrangement of noble gases are octet duplet, and this is the most stable electron arrangement of an atom.
• Atoms of other main group elements which is not octet tend to react with other atoms in various ways to achieve the octet.
• The tendency of an atom to achieve an octet arrangement of electrons in the outermost shell is called the octet rule.
• If the outermost shell is the first shell, then the maximum number of electrons is two, and the most stable electron arrangement will be duplet.
• A configuration of two electrons in the first shell, with no other shells occupied by electrons, is as stable as the octet electron arrangement and therefore is also said to obey the octet rule.

Important Notes Most of the elements (except noble gases) are chemically not stable. It is the aim of every atom to achieve the duplet or octet electron arrangement. This makes them very stable. It is only the valence electrons in the outermost shell involved in bonding. The electrons in the inner shells are not involved. The maximum number of electrons in the first shell is two. This is called a duplet. The maximum number in the second shell is eight. This is called an octet.

How Atoms Achieve Duplet or Octet Electron Arrangement?

• Atoms can achieve duplet or octet electron arrangement in 3 ways:

1. throw away the excess electron(s)
2. receiving electron(s) form other atom if they are lack of electron(s)
3. sharing electron

• 2 types of chemical bonds are commonly formed between atoms, namely

1. Ionic Bond
2. Covalent Bond

[edit] The Ionic Bond

• By releasing or receiving electron(s), the atoms will become ions and consequently form ionic bond between the ions.
• Ionic bonds are always form between metal and non-metal. For example, sodium (metal) react with chlorine (non-metal) will form an ionic bond between sodium ion and chloride ion.
• The compounds formed is called the ionic compound.
• Some time, an ionic bond is also called electrovalent bond
• The Covalent Bond

• By sharing electron(s), the atoms will form covalent bond between the atom and the molecule formed is call the covalent molecule.
• Covalent bond is always formed between non-metal with another non-metal.

Transition Elements And Its Properties

• Transition metal is a block of metallic elements in between Groups 2 and 13 in the Periodic Table.
• They are much less reactive than the alkali metals.
• They do not react as quickly with water or oxygen as alkali metal.

Some General Physical Characteristics

• All transition elements are metals. Therefore they have the all the physical properties of metal such as:

1. high melting point and boiling point
2. hard,
3. high density,
4. high electrical conductivity,
5. high tensile strength ,
6. shinny surfaces,
7. ductile
8. malleable ,

High Melting Point and Boiling Point

• The bond between atoms of metal is called metallic bond and usually it is a very strong bond.
• Thus all the transition metals have high melting points and boiling points.
• For example: iron melts at 1535°C and boils at 2750°C.
• Mercury is a transition metal, but with unusual low melting point ( − 39^oC).

Special Properties of Transition Elements

Form Coloured Compounds and Ions in Solute

• Transition elements tend toform coloured compounds either in solid form or dissolved in a solvent.
• Table below shows the colours of some aqueous solutions of ions of transition elements.

Ion	Colour
Fe2+	Light green
Fe3+	Brown
Ni2+	Green
Cr3+	Green
Mn2+	Pink
Cu2+	Blue
Co2+	Pink
MnO4-	Purple
CrO42-	Yellow
Cr2O72-	Orange

Catalytic Properties

• Transition elements or their compounds have catalytic properties.
• A catalyst is a substance that speeds up a reaction but it itself does not change chemically after a reaction.
• Many transition metals are used directly as catalysts in industrial chemical processes.
• Table below shows the uses of transition elements or their compounds as catalysts in industries .

Transition Elements or Its Compound	Uses
Platinum and rhodium	Used in the catalytic converters in car exhausts to reduce the emission of carbon monoxide and nitrogen
Platinum	Ostwald Process in the manufacture of nitric acid.
Nickel	Catalyst for 'hydrogenation' in the margarine industry.This process converts unsaturated vegetable oils into higher melting saturated fats.
Iron powder	Haber Process in the manufacture of ammonia. Iron catalyzes the reaction between nitrogen and hydrogen gas to produce ammonia.
Vanadium(V) oxide	Contact Process in the manufacture of sulphuric acid . In Contact Process, Vanadium(V) oxide catalyzes the oxidation of sulphur dioxide to sulphur trioxide.

Variable Oxidation State

• A transition metal can have a variable oxidation state, which means it can form more than one ion.
• Table below shows the ions of transition elements that have

for example, iron(II), Fe2+ and iron(III) Fe3+; copper(I), Cu+ and copper (II) Cu2+.

Forming Complex Ions

• Transition elements can form complex ions.
• A complex ion is a polyatomic ions (positive or negative) consisting of a central metal ion with other groups bonded to it.

The period

• Period is the horizontal rows of elements in the Periodic Table.
• The modern periodic table has 7 periods.
• The period number indicates the number of electron shell.
• Elements in the same period have same number of electron shells.
• The proton number of elements increases from left to right crossing the period.
• The number of electrons is also increases from left to right crossing the period.

Physical Change Across a Period

• There are a few major changes in physical properties across a period.
• For example,

1. the state of matter changes across the period.
2. The size of the molecule changes across the period, too.

State of Matter

• The diagram above shows that the state of matter of the elements in period 2 and period 3 change from solid to gas across the period (At room temperature).
• In period 2, the first 4 elements are solid whereas the last 4 elements are gases. The melting point of lithium is low while the melting point of boron and carbon is very high.
• In period 3, the first 6 elements are solid whereas the last 2 elements are gases. In fact, it is very difficult to find elements exist as liquid state at room temperature. The only one that I know is bromine.

Atomic Size

• As shown in the diagram above, the atomic size of elements in period 2 and period 3 decreases across a period.
• This is because

1. The number of proton in the nucleus increases whereas the number of electron shell remain unchanged across a period.
2. As the number of proton in the nucleus increases, the positive charge of the nucleus will also increases.
3. The negative charge of the valence electrons will also increases due to the increase of number of valence electrons across a period.
4. Thus, the attraction force between the nucleus and the valence electrons is getting stronger and stronger across a period.
5. This force will pulls the valence electrons closer to the nucleus and thus reduces the atomic radius.
6. Therefore the size of atom decreases across a period from left to right.

• Diagram below shows the change of the atomic size of the elements in all seven period. The trends can be conclude as below:

1. The size of atom increases down the group.
2. The size of atom decreases across a period from left to right.
3. For transition metal, the size of atom does not has obvious change across a period.

Chemical Properties Change Across a Period

• There is gradual chemical change across a period.

1. The acidity (or basicity) of the oxide of element changes across a period.
2. The metallic property, the electronegativity of element also change across a period.

Acidic Oxide or Basic Oxide

1. From left to right across a period the oxides change from alkaline/basic (with metals e.g. Na2O) to acidic (with non-metals e.g. SO2).

Metal, Metalloid and Non-metal

Metal and Non-metal

• As we go across a period from left to right, the elements change from metals to non-metals.
• Most of the known elements are metals.
• Only about 19 elements in the periodic table are non-metal.
• There are about 7 elements in the periodic table are classified as semi-metals.
• The metals in the periodic table are mainly found in the left hand columns (Groups 1 and 2) and in the central blocks of the transition elements.
• On the right hand side of the periodic table, there are 7 semi-metals (blue colour) form a staircase like pattern, act as a divider between metal and non-metal.
• The semi-metals are also called the metalloid.
• All the elements on the right hand side of the staircase are non-metals.

Uses of Metalloids

• The most widely used semi-metals are silicon and germanium.
• It is used to make diodes and transistors in electronic industry.
• Discussion of electronic can be found in SPM form 5 Physics.

Electronegativity

• Electronegativity is a measure of the potential of atoms to attract electrons to form negative ions.
• It is indicated bya number between 0 and 4.0. For example, the electronegativity of fluorine is 3.98 and the the electronegativity of sodium is 0.93.
• Metals have low electronegativities.
• Non-metals have high electronegativities.
• Electronegativities of the elements increase across a period with increasing proton number. This is because

1. as the proton number increases, the positive charge of the nucleus will increase accordingly.
2. this will increase the ability of the atom to attract electrons from the surrounding and thus increase the electronegativity of the atom.

• The table in the external link provided below shows the electronegativity of elements in periodic table. Nevertheless, the electronegativities of each elements are not important in SPM syllabus.

Introduction

• Group 17 elements are typical non-metals and also known as halogens.
• As shown in the diagram on the right, elements in this group (pink colour) are fluorine. chlorine, bromine, iodine and astatine.
• In nature, all halogens exist as diatomic molecules. They are written as F₂, Cl₂, Br₂, I₂and As₂.
• The word 'Halogens' means 'salt formers' because they can form salt easily with metals.
• For example, chlorine reacts with sodium forms sodium chloride. Sodium chloride is one of the most abundant salt found in the nature
• Most of the halogens exist in the nature as halide salts.
• Halide is the name given to the ion of halogens. Table below shows the corresponding halide of the halogen.

Halogen	Halide
Fluorine	Fluoride
Chlorine	CHloride
Bromine	Bromide
Iodine	Iodide

Physical Properties

• All group 17 elements are non-metals. Therefore they are heat and electricity insulator.
• Table below shows the electron arrangement and physical properties of group 17 elements.

Name	Proton Number	Electron arrangement	colour	Melting point	Boiling point
Fluorine	9	2.7	pale yellow gas	-220oC	-188oC 85K
Chlorine	17	2.8.7	Yellowish green gas	-102oC	-34oC
Bromine	35	2.8.18.7	dark red liquid, brown vapour	-7oC	59oC
Iodine	53	2.8.18.18.7	black solid, purple vapour	114oC	184oC
Astatine	85	2.8.18.32.18.7	black solid, dark vapour	302oC	380oC

Size of Atom and Density

• As shown in the diagram above, the atomic size of group 17 elements increases down the group.
• This is due to the increase of number of electron shell down the group.
• The density of group 17 elements is also increases down the group.
• This is because the rate of increment of the atomic mass is higher than the rate of increment of the volume.
• When solid of group 17 elements (Iodine and Astatine)they are brittle and crumbly.

Colour

• All halogens are coloured non-metallic elements.
• The colour of the halogen gets darker down the group.
• Diagram above shows the picture of the halogens.
• More high quality pictures of the elements can be found in [www.theodoregray.com www.theodoregray.com]. Click on the links provided below to access to the particular images.

Pictures of the Halogen: theodoregray.com: Fluorine theodoregray.com-wooden periodic table: Fluorine theodoregray.com: Chlorine theodoregray.com-wooden periodic table: Chlorine theodoregray.com: Bromine theodoregray.com-wooden periodic table: Bromine theodoregray.com: Iodine theodoregray.com-wooden periodic table: Iodine

Melting Point and Boiling Point

• As shown in the graph above, the melting points and boiling increase steadily down the group.
• The physical state t room temperature also change from gas to liquid and then to solid.
• This is because the intermolecular attractive force (van der Waals force) increase with increasing size of atom or molecule.

Chemical Properties

• Group 17 elements are very reactive non-metals.
• The atoms all have 7 valence electrons, makes them have very similar chemical properties.
• During chemical reaction, the atom gain one electron to form an ion with charge of -1.

• For example (as shown in the diagram above), the electron arrangement of fluorine is 2.7. In a chemical reaction, it will receive 1 electron from other atom and form fluoride ion, F^-. The diagram below shows the formation of chloride (Cl^-)ion when it receives 1 electron from other atom.
• The bromine and iodine will react in the same way to form bromide and iodide ions.
  

• In nature, the elements all exist as diatomic molecules, X₂ where X represents the halogen atom. For example F₂, Cl₂, Br₂, I₂ and At₂.
• The reactivity of group 1 decreases down the group.
• All group 17 elements are poisonous.
• Astatine is very radioactive.

• In the SPM chemistry syllabus, we discuss 3 reaction of group 17 elements:

1. Reaction with water.
2. Reaction with metal.
3. reaction with alkali solution, for example sodium hydroxide.

• The discussion of this 3 reaction are as follow.

React with water

Reaction of Chlorine with Water

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Observation:

• As shown in the diagram above, chlorine gas dissolves in water to form a pale yellow solution.
• When the solution is tested with blue litmus paper, the blue litmus paper turns red before it is bleached.
• This shown that the solution is acidic and contain bleaching agent.

Discussion:

• Chlorine gas dissolves in water produces hydrochloric acid (HCl) and hypochlorous(I) acid (HOCl).
• Hypochlorous(I) acid is a strong bleaching agent. It decolourises the colour of litmus paper.
• The equation of the reaction is shown below.

Chlorine + Water Hydrochloric acid + Hypochlorus acid

Reaction of Bromine with Water

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Observation:

• Bromine liquid dissolves slowly in water to form a yellowish-brown solution.
• When the solution is tested with blue litmus paper, the blue litmus paper turns red and then bleached slowly.

Discussion:

• Bromine liquid dissolves in water produces hydrobromic acid (HBr) and hypobromous(I) acid (HOBr).
• Hypobromous(I) acid is a weak bleaching agent.
• The equation of the reaction is shown below.

Bromine + Water Hydrobromic acid + Hypobromous(I) acid

Reaction of Iodine with Water

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Observation:

• Only a little iodine dissolves in water to form a yellowish solution.
• When the solution is tested with blue litmus paper, the blue litmus paper turns red but it is not bleached .

Discussion:

• Iodine solid slightly dissolves in water produces hydroiodic acid (HI) and hypoiodous(I) acid (HOI).
• Hypoiodous(I) acid has very weak bleaching characteristic..
• The equation of the reaction is shown below.

Iodine + Water Hydroiodic acid + Hypoiodous(I) acid

Chlorine, bromine and iodine are soluble in water to form an acidic solution. The solubility decreases down the group. Aqueous chlorine and bromine are bleaching agent. Aqueous iodine does not act as bleaching agent.

React with iron

• Halogen is a reactive non-metal. It form salt when react with metal.
• The experiments below show the reaction of chlorine, bromine and iodine with iron.
• Iron wool rather than iron piece is used to increase the rate of reaction.

Reaction of Chlorine with Iron

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Observation The iron wool burns vogorously with bright flame, forming a brown solid after reaction. Discussion The reaction of potassium with concentrated hydrochloric acid produces chlorine gas. The equation of the reaction is shown below. The chlorine gas flow into the combustion tube and react with iron wool when heated. Chlorine react with iron to form brown iron(III) chloride. Chlorine gas is a poisonous. Excess chlorine gas is absorbed by the soda lime (sodium hydroxide) so that it does not escape to the surrounding. Equation The equation of the reaction is as below: Chlorine + Iron iron(III) chloride

Reaction of Bromine with Iron

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Observation The iron wool glows brightly but less vigorously. A brown solid is formed. Discussion The bromine liquid evaporates to form bromine vapour when it is warmed. The bromine vapour flow into the combustion tube and react with iron wool when heated. Bromine react with iron to form brown iron(III) bromide. The reaction is less reactive compare with chlorine. Bromine gas is a poisonous. Excess bromine gas is absorbed by the soda lime (sodium hydroxide) so that it does not escape to the surrounding. Equation Bromine+ Iron iron(III) bromide

Reaction of Iodine with Iron

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Observation The iron wool glows slowly with dim light. A brown solid is formed after reaction. Discussion The iodine solid sublimates to form purplish iodine vapour when it is warmed. The iodine vapour flow into the combustion tube and react with iron wool when heated. Iodine react with iron to form brown iron(III) iodide. The reaction is slow compare with chlorine and bromine. Excess iodine vapour is absorbed by the soda lime (sodium hydroxide) so that it does not escape to the surrounding. Equation Iodine+ Iron iron(III) Iodide

React with Alkali Solution

Reaction of Chlorine with Sodium Hydroxide

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Observation:

• The greenish chlorine gas dissolves quickly in sodium hydroxide, forming a colourless solution.

Discussion:

• Chlorine gas react with sodium hydroxide to form salt of sodium chloride, sodium chlorate(I) and water.
• The equation of the reaction is shown below.

Chlorine + Sodium Hydroxide Sodium Chloride + Sodium Chlorate(I) + Water.

Reaction of Bromine with Sodium Hydroxide

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Observation:

• The reddish-brown liquid bromine dissolves in sodium hydroxide, forming a colourless solution.

Discussion:

• Bromine liquid react with sodium hydroxide to form salt of sodium bromide, sodium bromate(I) and water.

The reaction is less reactive compare with chlorine.

• The equation of the reaction is shown below.

Bromine+ Sodium Hydroxide Sodium Bromide + Sodium Bromate(I) + Water.

Reaction of Iodine with Sodium Hydroxide

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Observation:

• The black iodine crystals dissolve slowly in sodium hydroxide, forming a colourless solution.

Discussion:

• The reddish-brown liquid bromine dissolves in sodium hydroxide, forming a colourless solution.

Discussion:

• Solid iodine react slowly with sodium hydroxide to form salt of sodium iodide, sodium iodate(I) and water.

The reaction is leaet reactive among the three.

• The equation of the reaction is shown below.

Iodine + Sodium Hydroxide Sodium Iodide + Sodium Iodate(I) + Water.

Explaining the Reactivity Trend of the Group 17 Halogen

• The three experiments above shows that the reactiveness of halogens decreases down the group.
• This can be explained as below:

1. When a halogen atom reacts, it gains an electron to form a singly negative charged ion.
2. As we go down the group from F => Cl => Br => I, the size of the atom increases due to an extra filled electron shell.
3. The valence electrons are further and further from the nucleus, the attraction force between the electrons and the nucleus become weaker and weaker.
4. Therefore the ability of the atom to attract electron to fill the outermost shell reduces., which means the reactiveness of the atom reduces.

Safety Precaution

• Fluorine, chlorine and bromine gases are poisonous.
• Therefore all the experiments involving these gases should be carried out in a fume chamber.
• The experiments involve fluorine are nor done in school.
• This is because fluorine is so reactive that it will react with most of the substance it comes into contact with.
• It is very difficult to conduct experiments involving fluorine.

Introduction

• The Group I metals is called the Alkali Metals.
• This is because they form oxides and hydroxides that dissolve in water to give alkaline solutions.
• As shown in the diagram on the right, elements in this group are lithium, sodium, potassium, rubidium, caesium and fransium.
• They are the first element of a period, with one valence electron.
• This similarity (1 valence electron) makes them chemically behave in a similar manner.
• All alkali metals are very reactive. They must be stored in oil prevent reaction with oxygen or water vapour in air.

Physical Properties

Name	Proton number	Electron arrangement	melting point	boiling point	Density g/cm3
Lithium	3	2.1	180oC	1342oC	0.53
Sodium	11	2.8.1	98oC	883oC	0.97
Potassium	19	2.8.8.1	63oC	759oC	0.86
Rubidium	37	2.8.18.8.1	39oC	688oC	1.48
Caesium	55	2.8.18.18.8.1	29oC	671oC	1.87
Francium	87	2.8.18.32.18.8.1	27oC	677oC	> 1.87

• The table shows the physical properties and electron arrangement of Group 1 metals.
• All Group 1 metal exist as solid at room temperature.
• Since they are all metal, thus they have all the typical metallic properties, such as:

1. good conductors of heat
2. good conductors of electricity,
3. high boiling points,
4. shinny surface (but rapidly tarnished by air oxidation).

• Nevertheless, Group 1 metals also show some non-typical metallic properties, such as:

1. low melting points,
2. low density (first three float on water),
3. very soft (easily squashed, extremely malleable, can be cut by a knife).

Changes Down the Group

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Size of Atom

• Down the group, the size of atom increases.
• This is due to the increase of number of electron shells.
• Atom with more shells is bigger than atom with less shells.

Boiling Point and Melting Point

• The melting point and boiling point generally decrease down the group.
• All the atoms of Group 1 metals are bonded together by a force called metallic bond.
• The strength of metallic bond depends on the distance between the atoms.
• The nearer the atoms, the stronger the bond.
• Down the group, the size of the atoms increases, causing the distance of the atoms increases.
• As the distance between the atoms increases, the metallic bond between the atoms decreases.
• Therefore, less energy is needed to overcome the metallic bond during melting process.
• Consequently, the melting point of Group 1 metal decreases down the group.

Density

• The densities of Group 1 metals are low compare with the other metals.
• The densities of the first 3 elements (Lithium, Sodium and Potassium) are lower than water. Thus, they can float on the surface of water.
• Nevertheless, the density increases steadily down the group.
• Density of a substance is given by the equation .
• Down the group, both the mass and the volume increase, but increase of mass is faster than the volume, hence the density increases down the group.

Important trends down the group with increase in atomic number .. size of atoms increases the melting point and boiling point decrease the density increases. the hardness decreases.

Chemical Properties

• Group 1 metals are very reactive metals.
• They all show the same chemical properties.
• They can react with water and non-metal such as oxygen and chlorine to form a new compound.

Electron configuration

Lithium	2.1
Sodium	2.8.1
Potassium	2.8.8.1
Rubidium	2.8.18.1
Cesium	2.8.18.18.1
Francium	2.8.18.32.18.1

• Table above shows the electron arrangement of all the Group 1 metals.
• As we did mention before, all the atoms of Group 1 metal consist of 1 valence electron.
• When an alkali metal atoms reacts, it loses the valence electron to form a positively charged ion. Example

• They tend to react mainly with non-metals to form ionic compounds.
• We will discuss in detail the formation of ion and ionic compound in next chapter.
• In this chapter, we are going to discuss 3 example of reaction of Group 1 metal:

1. Reaction with water.
2. Reaction with chlorine gas.
3. Reaction with oxygen gas.

The Reaction of Alkali Metals with cold water

• Group 1 metal react vigorously with water.
• The video below shows the reaction of Lithium, Sodium, Potassium, Rubidium and Caesium with water.
• Click the play button to play the video or click anywhere on the player to link to the video page in youtube.

More Video Prentice Hall (Quicktime video)

Observation

Metal	Observation
Lithium	Lithium floats on the surface of the water with with 'fizzing' sound. Colourless gas is released around the metal. Lithium metal moves slowly on the surface of water. The gas released can be ignited. If the gas is collected in a test tube, and then a lighted wooden splinter is brought close to the mouth of the test tube, a "pop" sound is heard. The solution turn blue when it is tested with universal indicator.
Sodium	Sodium also floats on the surface of the water with with 'fizzing' sound. Colourless gas is released around the metal, as what happen to lithium. The lump of sodium moves swiftly on the surface of water. The gas released can be ignited. If the gas is collected in a test tube, and then a lighted wooden splinter is brought close to the mouth of the test tube, a "pop" sound is heard. The solution turn blue when it is tested with universal indicator.
Potassium	Sodium also floats on the surface of the water. It reacts violently with water. Colourless gas is released around the metal. The gas produceed is ignited by the heat release by the reaction itself. If the gas is collected in a test tube, and then a lighted wooden splinter is brought close to the mouth of the test tube, a "pop" sound is heard. The solution turn blue when it is tested with universal indicator.

Discussion

• The reaction with water is highly exothermic (release a lot heat), very fast and violent.
• If a lump of lithium, sodium or potassium is placed in cold water, the metal floats, move around the surface of the water and then dissolve in the water.
• This shows that lithium, sodium and potassium are less dense than water.
• The substance produced is soluble in water.

• The colourless flammable gas is hydrogen. It produces "pops" sound with lit splint.
• Lithium and sodium do not normally cause a flame but the potassium reaction is exothermic enough to ignite the hydrogen.
• Rubidium and caesium are explosive with water. Normally, your teacher does not do this experiment in the school lab.

• If universal indicator is added, it changes from green (pH 7) to purple (pH 13-14), showing that the products are alkali (solution of hydroxide).

• The more reactive the metal, the more vigorous the reaction.
• The reactivity of the metal increases down the group.

Equation of the reaction
Lithium + Water → Lithium Hydroxide + Hydrogen Gas

Sodium+ Water → SodiumHydroxide + Hydrogen Gas

Potassium + Water → Potassium Hydroxide + Hydrogen Gas

The Reaction of Alkali Metals with Non-metals

• Group 1 Alkali Metals react with non-metals to form colourless or white ionic compounds.
• These compounds dissolve in water to give colourless solutions.

Reaction with oxygen

• The diagram above shows that when hot alkali metal is put into a gas jar filled with oxygen gas, the alkali metal will burn with bright flame.
• They form white oxide powders after reaction.
• These oxides dissolve in water to form strongly alkaline metal hydroxide solutions with pH value 13-14.
• Lithium, sodium and potassium have similar chemical properties. All react with oxygen to produce white metal oxide.
• The reactivity increases down the group from lithium, sodium to potassium.
• Table below shows the observation of lithium, sodium and potassium when burn in oxygen gas.

Observation

Lithium	Lithium burns with red flame and produces white powder immediately after reaction. When the white powder is dissolved in water, it produces a solution which turned red litmus paper blue.
Sodium	Sodium burned with bright yellow flame, forming white powder immediately after reaction. When the white powder is dissolved in water, it produces a solution which turned red litmus paper blue.
Potassium	Potassium burned with very bright purplish flame, forming white powder immediately after reaction. When the white powder is dissolved in water, it produces a solution which turned red litmus paper blue.

Equation

Lithium+ Oxygen Lithium Oxide

Sodium + Oxygen Sodium Oxide

Potassium + Oxygen Potassium Oxide

Reaction with chlorine

• The diagram above shows that when hot alkali metal is put into a gas jar filled with chlorine gas, the alkali metal will burn with bright flame.
• All alkali metals react with chlorine gas to form white metal chlorides salt.
• Lithium, sodium and potassium have similar chemical properties.
• The metal chlorides salt formed is soluble in water to give a neutral solution of pH 7.
• The reactivity increases down the group from lithium, sodium to potassium.
• Table below shows the observation of lithium, sodium and potassium burn in chlorine gas.

Observation

Lithium	Lithium burned slowly with a reddish flame . A white solid is produced.
Sodium	Sodium burned brightly with a yellowish flame. A white solid is produced. (The reaction is shown in the youtube video below.)
Potassium	Potassium burned very brightly with a purplish flame . A white solid is produced.

Equation

Lithium + Chlorine Lithium Chloride

Sodium + Chlorine Sodium Chloride

Potassium + Chlorine Potassium Chloride

Others related video: www.youtube.com/watch?v=VBReOjo3ri8&eurl Prentice Hall (Quicktime video): Showing reaction of sodium with chlorine gas to form white sodium chloride salt.

Explaining the Reactivity Trend of the Group 1 Alkali Metals

• When an alkali metal atom reacts, it loses its valence electron to form a positively charged ion.
  

Example

• As we go down the group from one element down to the next, the atomic radius gets bigger due to an extra filled electron shell.
• The valence electron is further and further from the nucleus. Thus the attraction force between the nucleus and the valence electron become weaker and weaker.
• This causes the valence electron is easier to be released to form an ion when the atom takes part in a reaction.

Solubility of the Oxide, Hydroxide and Salt of Alkali Metal

• All the oxide and hydroxide of group 1 metal are soluble in water to form an alkali solution.
• All the salt (salt of chloride, nitrate, sulphate, carbonate....) of group 1 metal are soluble in water. The solution formed are neutral.

Safety Precaution

• Alkali metals are very reactive.
• Therefore it must be kept in paraffin oil to prevent them from reacting with oxygen and water vapour in tbe air.
• We must avoid to hold group 1 metals with bare hand because they may react with water on our hand.
• We must wear safety goggles and gloves during handling experiment involving group 1 metal.