Chemistry KS4 science GCSE/IGCSE O Level Revision Notes
The ALKALI METALS
Group 1 The Alkali Metals of the Periodic Table
Chemistry notes on the physical properties of lithium, sodium, potassium, rubidium, caesium (cesium) and francium, The chemical properties, chemical reactions with water, oxygen and chlorine – word equations & balanced equations and uses of 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.
Keywords–Links for this page: 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
Where are the Group 1 Alkali Metals in the Periodic Table?
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 the 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.
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.
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!
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.
Alkali metals form the first element of a period, with one outer electron, in any period from period 2 onwards. This outer electron similarity makes them behave in a chemically similar way and in a particularly reactive way.
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?
In what ways are the group 1 alkali metals like lithium, sodium and potassium typical metals?
In what ways are the group 1 alkali metals like lithium, sodium and potassium not typical metals?
Important trends down Group 1 Alkali Metals
With increase in atomic number (proton number), for the Alkali Metals ...
* For advanced AS level: 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.
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.
|more advanced data||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||181oC, 454K||1347oC, 1620K||0.53||0.157, 157|
|Na, sodium||11||2.8.1||98oC, 371K||883oC, 1156K||0.97||0.191, 191|
|K, potassium||19||220.127.116.11||64oC, 337K||774oC, 1047K||0.86||0.235, 235|
|Rb, rubidium||37||18.104.22.168.1||39oC, 312K||688oC, 961K||1.48||0.250, 250|
|Cs, caesium||55||22.214.171.124.8.1||29oC, 302K||679oC, 952K||1.87||0.272, 272|
|Fr, francium||87||126.96.36.199.18.8.1||27oC, 300K||677oC, 950K||approx. 2||~0.280, ~280|
1nm = 10–9m
1pm = 10–12m
nm x 1000 = pm
nm = pm/1000
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!
The Group 1 Alkali Metals are very reactive towards cold water.
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 and sodium do not normally cause a flame but the potassium reaction is exothermic enough to ignite the hydrogen.
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
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!
See also the GCSE/IGCSE Reactivity of Metals Notes for the reactivity of other metals compared to these Group 1 Alkali 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. For these reactions you can substitute Li (lithium) and K (potassium) to obtain the equations for other Group I Alkali Metals.
|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.
They form white solid powders which are ionic compounds eg (Na+)2O2–
sodium + oxygen ==> sodium oxide (word equation)
4Na + O2 ==> 2Na2O
4Na(s) + O2(g) ==> 2Na2O(s) (symbol equation with state symbols)
+ + ==>
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
Na2O + H2O ==> 2NaOH
Na2O(s) + H2O(l) ==> 2NaOH(aq) (symbol equation with state symbols)
|Reaction with chlorine||
What is formed when group 1 alkali metals like
lithium, sodium or potassium react with chlorine?
Alkali metals burn when heated in chlorine to form colourless ionic salt like compounds eg Na+Cl–. This is a very expensive way to make it! Its much cheaper to produce it by evaporating sea water.
sodium + chlorine ==> sodium chloride (word equation)
2Na + Cl2 ==> 2NaCl
2Na(s) + Cl2(g) ==> 2NaCl(s) (symbol equation with state symbols)
The sodium chloride is soluble in water to give a neutral solution pH 7, universal indicator is green.
|'normal molecular' and ionic formula, M = Li, Na, K etc.||
of the Alkali Metals (note the group
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.
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.
The oxides are white ionic solids, very soluble in water to form the metal hydroxide (see above).
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)
Colourless, soluble, neutral crystalline salts, 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)
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)
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) + CO2(g) ==> Na2CO3(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)
(much more details on pH, neutralisation, equations and salt preparations on "Acids, Bases and Salts")
|You will find more on theses sorts of equations on ....||
|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)
potassium nitrate ===> potassium nitrite + oxygen
2KNO3 ==> 2KNO2 + O2
2KNO3(s) ==> 2KNO2(s) + O2(g) (symbol equation with state symbols)
Li 2.1 ==> Li+ or +
Na 2.8.1 ==> Na+ or [2.8]+
K 188.8.131.52 ==> K+ or [2.8.8]+
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?
Explaining the Reactivity Trend of the Group 1 Alkali Metals
|Alkali Metals – Storylines – USES and ....|
Used as a heat transfer coolant in certain nuclear reactors because of its excellent heat conduction properties. The energized vapour is an orange–yellow and used in street lamps.
sodium Na+ salts
Common salt from sea water or underground deposits is sodium chloride, NaCl, and is the raw material for making sodium, hydrogen, chlorine and sodium chloride by electrolysis (see Group 7 Halogens notes).
'Soluble Aspirin' is the sodium salt of an organic acid. Salts of solid organic acids are usually more soluble than the acid itself.
Sodium hydrogen carbonate NaHCO3
Sodium hydrogencarbonate's old name is sodium bicarbonate, often referred to as 'bicarb', is used in baking soda, pharmaceutical products like indigestion tablets and fire extinguishers.
Sodium hydroxide NaOH
An industrially important alkali used in the manufacture of soaps, detergents, salts of acids (see Aspirin above), paper and ceramics.
PLEASE NOTE that these LINKS are for A Level Students ONLY
ADVANCED LEVEL INORGANIC CHEMISTRY Part 7 s–block Gp 1 Alkali Metals/Gp 2 Alkaline Earth Metals sub–index: 7.1 Introduction to s–block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals * 7.2 Group 1 data and graphs * 7.3 Group 2 data and graphs * 7.4 General trends down groups I & II and formulae *7.5 Oxygen reaction & oxides of s–block metals * 7.6 Water reaction & hydroxides of group 1/2 metals * 7.7 Acid reaction & salts of group1/2 metals * 7.8 chlorine reaction & halide of group I/II metals * 7.9 carbonates & hydrogen carbonates of s–block metals * 7.10 Solubility trends of groups 1/2 OH, NO3,SO4,CO3 compounds * 7.11 Thermal decomposition and stability of group 1 and group 2 carbonates & nitrates * 7.12 Uses of s–block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals and their compounds
equation keywords: 2Na(s) + 2H2O(l) ==> 2NaOH(aq) + H2(g) 4Na(s) + O2(g) ==> 2Na2O(s) Na2O(s) + H2O(l) ==> 2NaOH(aq) 2Na(s) + Cl2(g) ==> 2NaCl(s) NaOH (aq) + HCl(aq) ==> NaCl(aq) + H2O(l) NaOH(aq) + HNO3(aq) ==> NaNO3(aq) + H2O(l) 2NaOH(aq) + H2SO4(aq) ==> Na2SO4(aq) + 2H2O(l) 2NaOH(aq) + CO2(g) ==> Na2CO3(aq) + H2O(l) Na2CO3 + 2HCl(aq) ==> 2NaCl(aq) + H2O(l) + CO2(g) 2K(s) + 2H2O(l) ==> 2KOH(aq) + H2(g) 2Li(s) + 2H2O(l) ==> 2LiOH(aq) + H2(g) 2NaNO3(s) ==> 2NaNO2(s) + O2(g) 2KNO3(s) ==> 2KNO2(s) + O2(g) ) ==> 2NaCl(aq) + H2O(l) + CO2(g) 2K(s) + 2H2O(l) ==> 2KOH(aq) + H2(g) 2Li(s) + 2H2O(l) ==> 2LiOH(aq) + H2(g)
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