gcse Group 1 ALKALI METALS lithium sodium potassium physical properties balanced equations explaining reactivity trend chemical reactions uses KS4 science igcse O Level chemistry revision notes

GCSE Chemistry Notes: The Group 1 Alkali Metals Li Na K etc.

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Group 1 ALKALI METALS of the Periodic Table

Doc Brown's Chemistry KS4 science GCSE/IGCSE Chemistry Revision Notes

The physical and chemical properties of the Group 1 The Alkali Metals of the Periodic Table and their compound uses

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Sub-index for GCSE page on Alkali Metals

1. Where are the Group 1 Alkali Metals in the Periodic Table? and the significance of their electronic structure

2. General introduction to the Group 1 Alkali Metals

3. What are the important trends down Group 1 Alkali Metals

4. Selected data on the Group 1 Alkali Metals

5. The Reaction of Alkali Metals with cold water

6. Explaining the alkali metal reactivity trend

7. The Reaction of Alkali Metals with Non–metals

8. Some compounds of the alkali metals - formula patterns

9. Uses of alkali metals and their compounds

Alkali Metals GCSE/IGCSE multiple choice QUIZ – easier Foundation tier

Alkali Metals GCSE/IGCSE multiple choice QUIZ – harder Higher tier

Keywords–Links for this GCSE revision notes page on alkali metals

boiling points * chlorides * data on elements * density * electron arrangements

explaining reactivity trend * hydroxides * melting points * oxides * reaction with oxygen/chlorine

reaction with water * trends * typical properties * untypical properties * uses

BUT this page can act as a primer for the study of alkali metals lithium, sodium, potassium etc.

Chemistry notes on the physical properties of the alkali metals – lithium, sodium, potassium, rubidium, caesium (cesium) and francium, The chemical properties of alkali metals, chemical reactions of alkali metals with water, oxygen and chlorine – word equations & balanced equations for the reactions of alkali metals. The uses of alkali metals – the elements and compounds of the Group 1 Alkali Metals of the Periodic Table e.g. lithium, sodium & potassium etc. Also covered are explaining the group I alkali metal reactivity trend, uses of the alkali metals and alkali metal compounds. These revision notes on the alkali metals should prove useful for the new AQA, Edexcel and OCR GCSE (9–1) chemistry science courses.

Doc Brown's chemistry revision notes: basic school chemistry science GCSE chemistry, IGCSE chemistry, O level & ~US grades 8, 9 and 10 school science courses or equivalent for ~14-16 year old science students for national examinations in chemistry Revision notes on the physical and chemical properties of the group 1 alkali metals, chemical reactions of alkali metals, compounds of the alkali metals, explaining the reactivity trend of alkali metals, reactions with water, chlorine and oxygen, help when revising for AQA GCSE chemistry, Edexcel GCSE chemistry, OCR GCSE gateway science chemistry, OCR GCSE 21st century science chemistry GCSE 9-1 chemistry examinations.

TOP OF PAGE and sub-index for GCSE Alkali Metals page

1. Where are the Group 1 Alkali Metals in the Periodic Table?

and the significance of their electronic structure

Pd metals Part of the modern Periodic Table
Pd = period, Gp = group
metals => non–metals

Gp1 Gp2 Gp3 Gp4 Gp5 Gp6 Gp7 Gp0

1 ₁H_{Note
that hydrogen does not readily fit into any group} ₂He

2 ₃Li ₄Be _{atomic number}Chemical Symbol eg ₄Be ₅B ₆C ₇N ₈O ₉F ₁₀Ne

3 ₁₁Na ₁₂Mg ₁₃Al ₁₄Si ₁₅P ₁₆S ₁₇Cl ₁₈Ar

4 ₁₉K ₂₀Ca ₂₁Sc ₂₂Ti ₂₃V ₂₄Cr ₂₅Mn ₂₆Fe ₂₇Co ₂₈Ni ₂₉Cu ₃₀Zn ₃₁Ga ₃₂Ge ₃₃As ₃₄Se ₃₅Br ₃₆Kr

5 ₃₇Rb ₃₈Sr ₃₉Y ₄₀Zr ₄₁Nb ₄₂Mo ₄₃Tc ₄₄Ru ₄₅Rh ₄₆Pd ₄₇Ag ₄₈Cd ₄₉In ₅₀Sn ₅₁Sb ₅₂Te ₅₃I ₅₄Xe

6 ₅₅Cs ₅₆Ba Transition Metals ₈₁Tl ₈₂Pb ₈₃Bi ₈₄Po ₈₅At ₈₆Rn

₈₇Fr ₈₈Ra

Group 1 Alkali Metals * Reactive Metals of Groups 1 and 2 * Transition Metals
Post-transition metals - diagonally down and across Groups 3 to 6 (ignored semi-metal classification)

Non-metals - diagonally down and across Groups 3 to 7 * The very unreactive Group 0 noble gas non-metals

The Group I Alkali Metals are the first vertical column on the left of the Periodic Table, where you find most of the metallic elements. Therefore an Alkali Metal is the first element on a period from period 2 onwards.

Only the top portion of the periodic table is shown and remember metals tend to be on the left and the alkali metals form the first vertical column. Group 1 Alkali Metals also include the elements caesium/cesium (Cs) and radioactive francium (Fr) below rubidium, but are not shown in the above section of the periodic table.

The alkali metals are so named because they readily react with water to form an alkaline solution of the hydroxide e.g. sodium produced the well known alkali sodium hydroxide.

THINKING AHEAD: From a working knowledge of the position of Group 1 Alkali Metals in the periodic table you should be able to predict the number of outer electrons of a Group 1 Alkali Metal,

possible compound formulae of Group 1 Alkali Metals, reactions and symbol equations of Group 1 Alkali Metals and the probable reactivity of an Alkali Metal in Group 1 from its position in the periodic table,

Group 1 elements, on the left of the periodic table, and at the start of a period, have one outer electron and so you would expect them to be very reactive metals and readily lose the outer electron to form a singly charged positive ion.,

but the physical properties of group 1 alkali metals as well as their chemical reactions and reactivity trend are also explained.

It is the similarity in electron structure (1 electron in the outer shell) that makes the chemistry of group 1 alkali metals the same – group 1 chemistry!

Note: Using 0 to denote the Group number of Noble Gases is very historic now, since, compounds of xenon are known exhibiting a valency of 8. Because of the horizontal series of elements e.g. like the Sc to Zn block (10 elements), Groups 3 to 0 can also be numbered as Groups 13 to 18 to fit in with the actual number of vertical columns of elements. This can make things confusing, but there it is, classification is still in progress!

Electronic structure and reactivity of Group 1 Alkali Metals in the context of their position in the Periodic Table

On reaction metals readily form positive ions in compounds by losing electrons e.g.

sodium Na – e^– ===> Na⁺, magnesium Mg – 2e^– ===> Mg²⁺ or aluminium Al – 3e^– ==> Al³⁺

These are typical electron changes for metals in groups 1, 2 and 3 and the alkali metals are on the furthest left of the periodic table with on one outer electron and it is on the left-hand side of the periodic table you find the most metallic character.

Metallic elements on the left-hand side of the periodic table quite easily lose their few less strongly held outer electrons giving them a high reactivity in forming positive ions eg groups 1 and 2 readily lose 1 or 2 electrons respectively (as with Na and Mg above) to give an electronically stable noble gas structure with a full shell of outer electrons.

These outer electrons, particularly for the alkali metals, are also shielded by other inner electron shells and farthest from the nucleus and so are less strongly held and need less energy to give positive ions with a full outer shell of electrons.

Its energetically easier for an alkali metal to lose one electrons to form a stable positive ion than gain 7 electrons to form an unstable negative ion with a noble gas structure.

Alkali metals have to lose the least electrons to form a stable positive ion with a noble gas structure, this requires the least energy and makes the group 1 alkali metals the most reactive metals on the left-hand side periodic table.

These points and explanations are elaborated on by looking at the chemical reactions of alkali metals further down the page.

TOP OF PAGE and sub-index for GCSE Alkali Metals page

2. General introduction to the Group 1 Alkali Metals (see also data table below)

The Alkali Metals form Group 1 of The Periodic Table, and called so because they form oxides and hydroxides that dissolve in water to give alkaline solutions and the alkaline hydroxide is formed when the metals react with water.

metals

metals

metal ==> non–metal groups

Gp 1

Gp2

Gp3

Gp4

Gp5

Gp6

Gp7

Gp0

₃Li

2.1

a short section of the periodic table with group 1 electron arrangements

₁₁Na

2.8.1

₁₉K

2.8.8.1

Alkali metals form the first element of a period, with one outer electron, in any period from period 2 onwards. This outer electron similarity of the alkali metals makes them behave in a chemically similar (e.g. similar formulae) and in a particularly reactive way and is a modern pre–requisite of a set of elements belonging to the same group. BUT their similarity in physical properties and chemical reactions fits in well with Mendeleev's original conception of a group classification.

The single outer electron is easily lost, so alkali metals are very reactive and in their chemical reactions they form a singly charged positive ion e.g. Li⁺, Na⁺, K⁺ etc. and the ions formed have a very stable noble gas electron arrangement.

Some of their physical properties of Group 1 Alkali Metals are typical of metals and some are not so typical of metals.

Although Alkali Metals all have one outer electron and so similar physical and chemical properties, a characteristic of a periodic table group, BUT always watch out for trends down a group too. Overview of Periodic Table

Why are the group 1 alkali metals like lithium, sodium and potassium store under oil?

Alkali Metals are so reactive that they must be stored under oil in a well sealed glass container to minimise reaction with the oxygen or water vapour in air. When exposed to air, alkali metals rapidly tarnish as a layer of alkali metal oxide is formed.

Note: Oil is a water repellent and a physical barrier towards air and moisture BUT the alkali metals still gradually corrode as traces of oxygen and water diffuse through the oil to reach the very reactive lumps of alkali metal!

Alkali metals should never be touched with the naked hands or fingers, any sweat on you hand would react with the metal producing heat (burn from exothermic reaction), and skin irritation and blistering from the very strongly corrosive alkaline hydroxide that is formed (see reaction of alkali metals with water).

Apart deliberately doing the reaction with water as a demonstration experiment, all apparatus must be completely dry and solvents must be free of moisture in them.

The group 1 alkali metals are so reactive they can burst into flames spontaneously in air, particularly potassium, rubidium and caesium.

In what ways are the group 1 alkali metals like lithium, sodium and potassium typical metals?

Alkali metals are many typical metallic properties: e.g. good conductors of heat and electricity, high boiling points, silvery grey surface (but rapidly tarnished by air oxidation).

When an alkali metal atoms reacts, it loses an electron (oxidation) to form a singly positively charged ion eg Na ==> Na⁺ + e^–. In terms of electrons 2.8.1 ==> 2.8 and so forming a stable ion with a noble gas electron arrangement.

They tend to readily react with most non–metals to form ionic compounds which are usually soluble white solids.

In what ways are the group 1 alkali metals like lithium, sodium and potassium not typical metals?

Alkali metals have several untypical metallic properties: e.g.

low melting points, this means unusually for metals, the bonding between the metal atoms in the crystal is weak.

low density (the first three float on water – lithium, sodium and potassium),

very soft (easily squashed or cut with a knife, extremely malleable) and so they have little material strength.

TOP OF PAGE and sub-index for GCSE Alkali Metals page

3. What are the important trends down Group 1 Alkali Metals

With increase in atomic number (proton number), down the group for the Alkali Metals ...

==>==> ==> down group 1 alkali metals

... are the important trends to know about the group 1 alkali metals - trends in melting point, boiling point, reactivity, size of atom (atomic radius), density and physical strength as you go down the group with increase in atomic number (proton number) or relative atomic mass.

The melting point and boiling point generally decrease down Group 1 Alkali Metals (see data table below)

All alkali metals are very reactive and the element gets more reactive down Group 1 Alkali Metals with increase in atomic number (explanation), though they all behave in a similar manner – all in the same group!

The atoms get bigger down Group 1 Alkali Metals (as more electron shells are added, see data table below)

Generally the density increases down Group 1 Alkali Metals (see data table below), although the atom gets bigger, there is a greater proportional increase in the atomic mass (see data table).

Generally the physical hardness decreases down Group 1 Alkali Metals, this suggests the metallic bonding gets weaker down the group.

* Note for advanced A level chemistry students: The bonding in metals involves the attraction between free negative electrons moving between positively ionised metal atoms (M⁺ ions). As the atomic radius increases the charges (positive nucleus and delocalised electrons) are further apart and the electrical attractive force is reduced. This weaker bonding results in a weaker–softer structure with a lower melting/boiling point.

A Level Chemistry AS–A2 Notes on s–block Gp 1 Alkali Metals & Gp2 Alkaline Earth Metals

Alkali Metal flame colours

There are element/compound identification details of this and other metal ion tests on the Chemical Tests page (use the alphabetical list at the top).

When heated strongly in a flame, the Alkali metals or their compounds give bright colours.

Simple method for a flame test: The metal salt or other compound is mixed with concentrated hydrochloric acid and a sample of the mixture is heated strongly in a bunsen flame on the end of a cleaned nichrome wire (or platinum if you can afford it!). Before doing the test the nichrome/platinum wire should be cleaned in conc. hydrochloric acid and heated in the hottest part of the flame to make sure there is no contaminating flame colours.

Lithium – red/crimson (carmine–red)

sodium – yellow/orange

potassium – lilac/purple

rubidium red, caesium – blue

and these flame colours of alkali metals can be used as a simple identification flame test for a specific alkali metal.

Revision notes on the physical and chemical properties of the group 1 alkali metals, chemical reactions of alkali metals, compounds of the alkali metals, explaining the reactivity trend of alkali metals, reactions with water, chlorine and oxygen, help when revising for AQA GCSE chemistry, Edexcel GCSE chemistry, OCR GCSE gateway science chemistry, OCR GCSE 21st century science chemistry GCSE 9-1 chemistry examinations.

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more advanced data	4. Selected data on the Group 1 Alkali Metals
Chemical symbol, name of alkali metal	Atomic number of alkali metal	Electron arrangement in shells 1, 2, 3 etc. of alkali metal	melting point ^oC, K of alkali metal	boiling point ^oC, K of alkali metal	Density g/cm³ of alkali metal	atomic radius in nm (nanometre) & pm (picometre)
Li, lithium	3	2.1	181^oC, 454K	1347^oC, 1620K	0.53	0.157, 157
Na, sodium	11	2.8.1	98^oC, 371K	883^oC, 1156K	0.97	0.191, 191
K, potassium	19	2.8.8.1	64^oC, 337K	774^oC, 1047K	0.86	0.235, 235
Rb, rubidium	37	2.8.18.8.1	39^oC, 312K	688^oC, 961K	1.48	0.250, 250
Cs, caesium	55	2.8.18.18.8.1	29^oC, 302K	679^oC, 952K	1.87	0.272, 272
Fr, francium	87	2.8.18.32.18.8.1	27^oC, 300K	677^oC, 950K	approx. 2	~0.280, ~280
GROUP 1 ALKALI METALS **************************	Proton number of group 1 alkali metals ***************	All group 1 alkali metals have one electron in the outer shell *****************************	The melting points of group 1 alkali metals decrease down the group. **********************	Generally, the boiling points of group 1 alkali metals decrease down the group **********************	Alkali metals have low densities, untypical of metals. **************	Atomic radii increase down the group with each additional electron shell. **************************
Note: For atomic radii: 1nm = 10^–9m, 1pm = 10^–12m, nm x 1000 = pm, nm = pm/1000 Atomic radii always increase down a group with increase in atomic number because extra electron shells are successively added.

TOP OF PAGE and sub-index for GCSE Alkali Metals page

What is formed when group 1 alkali metals like lithium, sodium or potassium react with water? What do you see when the reaction takes place? Observations!

Elements – metals on the far left–hand side of the periodic table quite easily lose their outer electrons giving them a high reactivity in forming positive ions. The group 1 alkali metals are the most reactive metal group of the periodic table and only have to lose their one outer electron to form a very stable mono–positive ion with a noble gas electron arrangement.

The Group 1 Alkali Metals are very reactive towards cold water producing hydrogen gas and an alkaline solution of the alkali metal hydroxide.

The reaction of alkali metals with water is very exothermic, fast and violent.

If a lump of the alkali metals lithium, sodium or potassium is placed in cold water, the metal floats, it may melt and move around the surface of the water with 'fizzing'.

Note that these alkali metals float on water because of their low density.

If universal indicator is added, it changes from green (pH 7) to purple (pH 13–14), showing an alkaline metal hydroxide was formed.

The formation of an alkali with water is why they are called Alkali Metals.

The colourless gas hydrogen is also given off and pops with lit splint – but this is not the best of experiments to collect it from!

The more reactive the alkali metal, the more vigorous the reaction.

Lithium reacts moderately fast. Sodium reacts faster. Potassium reacts very fast.

Rubidium is even faster and caesium is pretty explosive on contact with water.

The clearly observed reactivity trend is: Li > Na > K > Rb > Cs

You would NOT do the experiment with francium unless you want to explosively coat the laboratory walls with radioactive isotopes!

Lithium and sodium do not normally cause a flame but the potassium reaction is more rapid exothermic, enough to ignite the hydrogen.

hydrogen + oxygen ==> water (word equation)

2H₂ + O₂ ===> 2H₂O

2H₂(g) + O₂(g) ===> 2H₂O(l) (symbol equation with state symbols)

The hydrogen flame turns lilac due to hot potassium atoms in it.

Note 1: In flame tests hot lithium atoms give a crimson colour and sodium a yellow colour (more details).

Note 2: State symbols in equations: (g) = gas, (l) = liquid, (s) = solid, (aq) = aqueous solution in water

The alkali metals Rubidium, caesium and francium are very explosive with water.

Down group 1 the reaction gets faster and more violent as the alkali metal gets more reactive

i.e. Li < Na < K < Rb < Cs < Fr or francium > caesium > rubidium > potassium > sodium > lithium

The reaction equation for sodium is in words and symbols ...

sodium + water ==> sodium hydroxide + hydrogen    (word equation)

2Na + 2H₂O ==> 2NaOH + H₂

2Na(s) + 2H₂O(l) ==> 2NaOH(aq) + H₂(g)   (symbol equation with state symbols)

and the equations are similar for any of the other Alkali Metals since they are in the same group of the Periodic Table, they behave chemically in the same way. On the basis of what you have seen for Li, Na and K, you should therefore be able to predict the reactions and relative reactivity of Rb, Cs and Fr.

i.e. just substitute Li (lithium), K (potassium), Rb (rubidium), Cs (rubidium) or Fr (francium) for Na.

Its handy to know the 'pattern' of both the word equations and symbol equations and it fits in with the idea, an important concept, that elements in the same group of vertical column in the periodic table, all tend to behave chemically in the same way ...

... this is what the Periodic Table is all about!.

e.g. by learning how to write and balance the symbol equation for the reaction between sodium and water, all you have to do is get the symbol of another group 1 alkali metal from your periodic table and deduce the equation i.e. swap Na for a K etc.

potassium + water ==> potassium hydroxide + hydrogen    (word equation)

2K + 2H₂O ==> 2KOH + H₂

2K(s) + 2H₂O(l) ==> 2KOH(aq) + H₂(g)   (symbol equation with state symbols)

and

lithium + water ==> lithium hydroxide + hydrogen   (word equation)

2Li + 2H₂O ==> 2LiOH + H₂

2Li(s) + 2H₂O(l) ==> 2LiOH(aq) + H₂(g)   (symbol equation with state symbols)

You can predict that the continued reactivity order after Li, Na, K will be K < Rb < Cs < Fr

You can write similar word or symbol equations for rubidium, caesium and francium by replacing Li, Na or K with Rb, Cs or Fr - similar formula and similar balanced equations - all fully predictable without doing any extra experiments!

Theoretically the alkali metal Francium is the most reactive alkali metal and therefore the most explosive metal when in contact with water, however, it is also very radioactive and so the experiment is highly unlikely to be performed!

The explanation of the alkali metals reactivity and reactivity trend is explained and discussed in the next section

See also the Reactivity of Metals Notes for the reactivity of other metals compared to these Group 1 Alkali Metals.

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6. Why are alkali metals so reactive? AND

Why do Group 1 Alkali Metals get more reactive down the group with increase in atomic/proton number? How do we explain the group 1 alkali metal reactivity trend?

metals

metals

metal ==> non–metal groups

Gp 1

Gp2

Gp3

Gp4

Gp5

Gp6

Gp7

Gp8

₃Li

2.1

a short section of the periodic table with group 1 electron arrangements

₁₁Na

2.8.1

₁₉K

2.8.8.1

Representations of the electronic changes that accompany the reaction of a group 1 alkali metal.

The atom loses its outer electron to form a very stable ion with the electron configuration of a noble gas.

==> + e^–

Li or [2.1] ==> Li⁺ or [2]⁺ + e^–

==> + e^–

Na or [2.8.1] ==> Na⁺ or [2.8]⁺ + e^–

==> + e^–

K or [2.8.8.1] ==> K⁺ or [2.8.8]⁺ + e^–

THEORY (c) doc b

Explaining the Reactivity Trend of the Group 1 Alkali Metals with water and acids

The electronic structure of the alkali metals is the basis for explaining their high reactivity and reactivity order.
When an alkali metal atom reacts, it readily loses its only outer electron to form a singly positively charged ion
- e.g. Na ==> Na⁺ + e^– (in terms of electrons 2.8.1 ==> 2.8)
- and so forming a very stable ion with a noble gas electron arrangement, (see diagrams on left)
- and this is why they are so reactive,
- the formation of a positive ion by electron loss is an example of oxidation.
AND as you go down the group from one element down to the next
.. Li .. Na .. K .. Rb .. Cs .. Fr the atomic radius gets bigger due to the addition of an extra filled inner electron shell,
the further a negative electron in the outer shell is from the positive nucleus, the less strongly it is held, hence the more easily lost from an atom to form a positive ion,
so, down the alkali metals group, the outer electron is further and further from the nucleus and is also shielded by the extra full shell of negative electron charge,
these combined effects mean that down the alkali metal group the outer electron is less and less strongly held by the positive nucleus as the attractive force is decreased, and so ....
this means the outer electron is more easily lost in the reaction (e.g. with water), the M⁺ ion is more easily formed, and so the alkali metal is more reactive as you go down the group.
See also 'The Reactivity Series of Metals'

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7. The Reaction of Alkali Metals with Non–metals

Group 1 Alkali Metals react with non–metallic elements to form colourless or white ionic compounds

These compounds dissolve in water to give colourless solutions.

For these reactions you can often substitute Li (lithium), K (potassium), Rb (rubidium), Cs (caesium/cesium) to obtain the equations for other Group I Alkali Metals.

The reason why they are ionically bonded compounds is that alkali metals readily form positive ions e.g. for sodium

===> + e^–

Na or [2.8.1] ==> Na⁺ or [2.8]⁺ + e^–

AND non-metals like oxygen and chlorine also readily form negative ions.

Reaction with oxygen

What is formed when group 1 alkali metals like lithium, sodium or potassium react with oxygen (air)?

Alkali metals burn when heated in oxygen or air – very exothermic reactions.

Unfortunately its a little bit complicated because the principal product, of several possible products, depends on the alkali metal.

They form white solid powders which are ionic compounds eg (Na⁺)₂O^2–

sodium + oxygen ==> sodium oxides (word equation)

4Na + O₂ ==> 2Na₂O

4Na(s) + O₂(g) ==> 2Na₂O(s) (symbol equation with state symbols)

Another oxide is also formed: 2Na + O₂ ===> Na₂O₂ (sodium peroxide)

he diagram below shows the ionic bonding of sodium oxide

+ + ==>

These oxides dissolve in water to form strongly alkaline metal hydroxide solutions, pH 13–14, so universal indicator turns from green to blue.

eg sodium oxide + water ==> sodium hydroxide (an alkali)

Na₂O + H₂O ===> 2NaOH

Na₂O(s) + H₂O(l) ===> 2NaOH(aq) (symbol equation with state symbols)

Similarly lithium forms lithium oxide: 4Li(s) + O₂(g) ===> 2Li₂O(s)

but potassium, like sodium, forms several potassium oxides including:

2K + O₂ ===> K₂O₂ and K + O₂ ===> KO₂

(this reaction is more complicated than for lithium, but I wouldn't worry about it at GCSE level!)

and they all dissolve in water to form hydroxides which are very alkaline, pH ~13/14, litmus turns blue, universal indicator turns blue–violet.

Reaction with chlorine and other halogens

What is formed when group 1 alkali metals like lithium, sodium or potassium react with chlorine?

Alkali metals burn vigorously when heated in chlorine to form colourless ionic salt like compounds eg Na⁺Cl^–. This is a very exothermic reaction AND expensive way to make it! Its much cheaper to produce it by evaporating sea water.

You can do the experiment by heating a small amount of the alkali metal in a deflagrating spoon and plunging into a gas jar of previously prepared chlorine when the metal will burn quite vigorously.

sodium + chlorine ==> sodium chloride (word equation)

2Na + Cl₂ ==> 2NaCl

2Na(s) + Cl₂(g) ==> 2NaCl(s) (symbol equation with state symbols)

The sodium chloride is an ionic compound, soluble in water to give a neutral solution pH 7, universal indicator is green.

Similarly lithium forms lithium chloride: 2Li(s) + Cl₂(g) ==> 2LiCl(s)

and potassium forms potassium chloride: 2K(s) + Cl₂(g) ==> 2KCl(s)

All the alkali metals react with all of the halogens to produce white crystalline solids of the ionic compound.

e.g sodium reacts with fluorine gas to give sodium fluoride.

2Na(s) + F₂(g) ===> 2NaF(s)

Since the charge on the group 1 metal ions is +1, and the charge on halide ions is -1, its easy to predict the formula of any ionic compound formed between an alkali metal and a halogen i.e. a 1 : 1 ratio.

Pd metals Part of the modern Periodic Table
Pd = period, Gp = group
metals => non–metals

Gp1 Gp2 Gp3 Gp4 Gp5 Gp6 Gp7 Gp0

1
₁H_{Note
that H does not readily fit into any group}
₂He

2 ₃Li ₄Be _{atomic number}Chemical Symbol eg ₄Be
Group 1 Alkali Metals and Group 7 Halogens
₅B ₆C ₇N ₈O ₉F ₁₀Ne

3 ₁₁Na ₁₂Mg ₁₃Al ₁₄Si ₁₅P ₁₆S ₁₇Cl ₁₈Ar

4 ₁₉K ₂₀Ca ₂₁Sc ₂₂Ti ₂₃V ₂₄Cr ₂₅Mn ₂₆Fe ₂₇Co ₂₈Ni ₂₉Cu ₃₀Zn ₃₁Ga ₃₂Ge ₃₃As ₃₄Se ₃₅Br ₃₆Kr

5 ₃₇Rb ₃₈Sr ₃₉Y ₄₀Zr ₄₁Nb ₄₂Mo ₄₃Tc ₄₄Ru ₄₅Rh ₄₆Pd ₄₇Ag ₄₈Cd ₄₉In ₅₀Sn ₅₁Sb ₅₂Te ₅₃I ₅₄Xe

6 ₅₅Cs ₅₆Ba Transition Metals ₈₁Tl ₈₂Pb ₈₃Bi ₈₄Po ₈₅At ₈₆Rn

So, you can deduce the following table of all the possible ionic compounds formed from the Group 1 Alkali Metals and the non-metal Group 7 Halogen elements.

Gp1\7 F Cl Br I

Li LiF LiCl LiBr LiI

Na NaF NaCl NaBr NaI

K KF KCl KBr KI

Rb Rb RbCl RbBr RbI

Cs CsF CsCl CsBr CsI

'normal molecular' and ionic formula, M = Li, Na, K etc.	8. Some compounds of the alkali metals - more on formula patterns Alkali metals are so reactive in readily forming a singly charged positive ion, they usually form ionic compounds, they lose an electron and are NOT interested in sharing it to form a covalent bond! e.g. the formation of the ionic compound sodium chloride (see above) ONE atom combines with ONE atom to form Alkali metal compounds are usually white solids or colourless crystalline compounds.
hydroxides MOH, M⁺OH^–	The hydroxides are white ionic solids which very soluble in water to form strongly alkaline solutions (pH 13–14). See below for salt formation from hydroxides.
oxides, M₂O (M⁺)₂O^2–	The oxides are white ionic solids, very soluble in water to form the metal hydroxide (see above).
chlorides MCl M⁺Cl^–	The chlorides are colourless crystalline solids. They soluble in water to give a neutral solution pH 7, universal indicator is green. They are typical ionic solids with high melting points due to the strong attractive forces between ions (ionic bonding details). This solution in water consists of sodium Na⁺ and chloride Cl^– ions and can be electrolysed to make chlorine, hydrogen and sodium hydroxide. Formed by neutralising the alkaline oxide or hydroxide with acids (more on Acids, Bases and Salts). e.g. word equation and symbol equations sodium hydroxide + hydrochloric acid ===> sodium chloride + water NaOH + HCl ===> NaCl + H₂O NaOH(aq) + HCl(aq) ===> NaCl(aq) + H₂O(l) (symbol equation with state symbols)
nitrates MNO₃ M⁺NO₃^–	Colourless, soluble, neutral crystalline salts, are formed by neutralising the alkaline oxide or hydroxide with nitric acid. e.g. word equation and symbol equations sodium hydroxide + nitric acid ===> sodium nitrate + water NaOH + HNO₃ ===> NaNO₃ + H₂O NaOH(aq) + HNO₃(aq) ===> NaNO₃(aq) + H₂O(l) (symbol equation with state symbols)
sulphates M₂SO₄ (M⁺)₂SO₄^2–	Colourless, soluble, neutral crystalline salts, formed by neutralising the alkaline oxide or hydroxide with sulphuric acid. e.g. word equation and symbol equations sodium hydroxide + sulphuric acid ===> sodium sulphate + water 2NaOH + H₂SO₄ ==> Na₂SO₄ + 2H₂O 2NaOH(aq) + H₂SO₄(aq) ==> Na₂SO₄(aq) + 2H₂O(l) (symbol equation with state symbols)
carbonates M₂CO₃ (M⁺)₂CO₃^2–	White, soluble, weakly alkaline solids formed by reacting the hydroxide with carbon dioxide gas e.g. the formation of sodium carbonate (+ water) e.g. word equation and symbol equations sodium hydroxide + carbon dioxide ===> sodium carbonate + water 2NaOH + CO₂ ===> Na₂CO₃ + H₂O 2NaOH_(aq) + CO_2(g) ===> Na₂CO_3(aq) + H₂O_(l) (symbol equation with state symbols) Alkali metal carbonates form salts with acids. e.g. sodium carbonate + hydrochloric acid ===> sodium chloride + water + carbon dioxide Na₂CO₃+ 2HCl ===> 2NaCl + H₂O + CO₂ Na₂CO₃(s) + 2HCl(aq) ===> 2NaCl(aq) + H₂O(l) + CO₂(g) (symbol equation with state symbols) (more details on pH, neutralisation, equations and salt preparations on "Acids, Bases and Salts")
You will find more on theses sorts of equations on ....	GCSE Notes on Acids, Bases, Salts and pH scale Equation Question on Acid reactions – word & symbol equations with answers supplied
Other reactions involving alkali metals or alkali metal compounds	IGCSE chemistry may need the effect of heat on alkali metal nitrates. When strongly heated the nitrates of sodium and potassium evolve oxygen gas (ignites glowing splint) and leaving a white residue of the nitrite salt i.e. sodium nitrate ===> sodium nitrite + oxygen (word equation) 2NaNO₃ ===> 2NaNO₂ + O₂ 2NaNO₃(s) ===> 2NaNO₂(s) + O₂(g) (symbol equation with state symbols) and potassium nitrate ===> potassium nitrite + oxygen 2KNO₃ ===> 2KNO₂ + O₂ 2KNO₃(s) ===> 2KNO₂(s) + O₂(g) (symbol equation with state symbols)

Equation Question on Acid reactions – word & symbol equations with answers supplied