S-Block Elements ALKALI METALS
Chapter summary • Introduction • Characteristic properties of the s-block elements . • Variation in properties of the s-block elementsof the First Group(Alkali Metals). • Physical Properties
IA IIA Li Be Mg Na Ca K Sr Rb Cs Ba Ra Fr Members of the s-Block Elements IA Alkali metals IIA Alkaline Earth metals
Characteristic properties of s-block elements • Metallic character • Low electronegativity • Basic oxides, hydroxides • Ionic bond with fixed oxidation states • Characteristic flame colours • Weak tendency to from complex
Metallic character • High tendency to lose e- to form positive ions • Metallic character increases down both groups
Electronegativity • Low nuclear attraction for outer electrons • Highly electropositive • Small electronegativity
Characteristic flame colours Na+ Cl- (g) Na(g) + Cl(g) Na(g) Na* (g) [Ne]3s1 [Ne]3p1 Na*(g) Na(g) + h (589nm, yellow)
HCl(aq) sample Flame test Li deep red Na yellow K lilac Rb bluish red Cs blue Ca brick red Sr blood red Ba apple green
Variation in properties of elements • Atomic radii • Ionization enthalpies • Hydration enthalpies • Melting points • Reactions with oxygen, water, hydrogen and chlorine
Atomic radii (nm) Fr Ra Li Be
1st I.E. 2000 600 Be+ Li Na 500 1500 2nd IE K 400 Rb Ca+ 1000 Cs Ba+ 300 Be Ca 500 Ba 1st IE Ionization Enthalpy
Ionization Enthalpy • Group I • Have generally low 1st I.E. as it is well shielded from the nucleus by inner shells. • 2. Removal of a 2nd electron is much more difficult because it involves the removal of inner shell electron. • 3. I.E. decreases as the group is descended. • As atomic radius increases, the outer e is further away from the well-shielded nucleus.
-600 -300 M+ Li+ Na+ K+ Rb+ Cs+ Hydration Enthalpy M+(g) + aqueous M+(aq) + heat
-2250 -600 -2000 -1750 -300 -1500 Li+ Na+ K+ Rb+ Cs+ Be2+ Mg2+ Ca2+ Sr2+ Ba2+ Hydration Enthalpy
Hydration Enthalpy • General trends: • On going down both groups, hydration enthalpy • decreases. • (As the ions get larger, the charge density of the • ions decreases, the electrostatic attraction between • ions and water molecules gets smaller.) • Group 2 ions have hydration enthalpies higher • than group 1. • ( Group 2 cations are doubly charged and have • smaller sizes)
Ionization Energy • Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. • First ionization energy is that energy required to remove first electron. • Second ionization energy is that energy required to remove second electron, etc.
Trends in First Ionization Energies • As one goes down a column, less energy is required to remove the first electron. • For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.
Ionization Enthalpy of Alkali Metals • ionization enthalpy decreases down the group from Li to cs because as we move down a group the number of valence electrons goes increasing separating the electrons away from the nucleus ,there is an increasing shielding of the nuclear charge by the inner shell electrons and thus the removal of electrons requires less energy as we move down.
Physical Properties • Silvery White • Soft & Light Metals • Due to large size , elements have low density • Low Melting & Boling Points indicates weak bonding due to presence of only 1 valence electron .
AlkaliMetals • In order of increasing atomic number the alkali metals are: • Lithium • Sodium • Potassium • Rubidium • Caesium • Francium Increasing atomic number
PropertiesofAlkali Metals All alkali (group 1) metals react violently with water, forming Hydrogen gas and Hydroxides (pH above 7): • Alkali metals are: • Metals found in group 1 of the periodic table. • Soft when cut (compared to other metals). • Metals with low melting points and densities. • Powerful reducing agent and form univalent compounds. • Metals which tarnish in air.
Oxides • On cumbustion in excess of air, • alkali metals form • Oxides- LiO2 • 2. Peroxides- Li2O2, NaO2 • 3. Superoxides-K2O2 ,Cs2O2 RbO2
Hydroxides The oxides of the alkali metal are easily hydrolysed by addition of water. E.g.:- LiO2 + H2O → LiOH
Halides The alkali metals combine directly with halogens under appropriate conditions forming halides of general formula MX. These halides can also be prepared by the action of aqueous halogen acids (HX) on metals oxides, hydroxides or carbonate. All these halides are colourless, high melting crystalline solids having high negative enthalpies of formation.
Examples of halides M2O + 2HX → 2MX + H2O MOH + HX → MX + H2O M2CO3 + 2HX → 2MX + CO2 + H2O (M = Li, Na, K, Rb or Cs) (X = F, Cl, Br or I)
Salts of oxo-acids Since the alkali metals are highly electropositive, therefore their hydroxides are very strong bases and hence they form salts with all oxoacids. ( H2CO3, H3PO4, H2SO4, HNO3, HNO2 etc) . They are generally soluble in water and stable towards heat.
The carbonates (M2CO3) of alkali metals are remarkably stable upto 1273 K, above which they first melt and then eventually decompose to form oxides. Li2CO3 , however is considerably less stable and decomposes readily. Δ Li2CO3→ Li2O + CO2 This is presumably due to large size difference between Li+ and CO2-3 which makes the crystal lattice unstable.
SODIUM CHLORIDE The most abundant source of sodium chloride is sea water. Crude sodium chloride, generally obtained by crystallisation of brine solution, contains sodium sulphate, calcium sulphate, calcium chloride and magnesium chloride as impurities. Calcium chloride,CaCl2, and magnesium chloride MgCl2 are impurities because they are deliquescent(absorb moisture from the atmosphere).
To obtain pure sodium chloride– Crude salt is dissolved in minimum amount of water and filtered to remove insoluble impurities. The solution is then saturated with hydrogen chloride gas. Crystals of pure sodium separate out. Calcium and magnesium chloride, being more soluble than sodium chloride, remains in solution.
USES OF NaCl − It is used as a common salt or table salt for domestic purpose. It is used for the preparation of Na2O2 , NaOH and Na2CO3.
SODIUM HYDROXIDE(CAUSTIC SODA), NaOH Sodium hydroxide is generally prepared commercially by the electrolysis of sodium chloride in Castner −Kellner cell. A brine soln. is electrolysed using a mercury cathode and a carbon anode. Sodium metal discharged at the cathode combines with mercury to form sodium amalgam. Chlorine gas is evolved at the anode.
NaCl → Na+ + Cl¯AT ANODE: Cl ¯ ─ e¯ → ClCl + Cl → Cl2AT CATHODE: Na+ + e¯ →Na+Na + Hg → NaHgAmalgamThe amalgam is treated with water to give sodium hydroxide, mercury and hydrogen gas.2NaHg + 2H2O→ 2NaOH + 2Hg + H2
PROPERTIES-1. NaOH is a white, translucent solid and it melts at 591K.2. it is readily soluble in water to give alkaline solution.3. Crystals of NaOH are deliquescent. It reacts with the CO2 in the atmosphere to form Na2CO3. USES- The manufacture of soap, paper and no. of chemicals. In petroleum refining In textile industry for mercerising cotton fabrics As laboratory reagent For preparation of pure oils and fats
SODIUM HYDROGENCARBONATE(BAKING SODA), NaHCO3 Sodium hydrogencarbonate is known as baking soda because it decomposes on heating to generate bubbles of carbon dioxide. It is made by saturating a solution of sodium carbonate with carbon dioxide. The white crystalline powder of sodium hydrogencarbonate , being less soluble, gets separated out. Na2CO3 + H2O + CO2 → 2NaHCO3 NaHCO3 is a mild antiseptic for skin infections. It is used in fire extinguishers.
BIOLOGICAL IMPORTANCE OF SODIUM AND POTASSIUM A typical 70 kg man contains about 90g of Na and 170g of K compared with only 5g of iron and 0.06g of copper. Potassium ions are present in higher concentration inside the cells than sodium ions and they are present outside the cell in blood plasma. Because of large concentration gradient inside and outside the cells, the transport of sodium ion into the cells is favoured. To pump out these ions again from the cell to maintain concentration gradient large driving force is carried out. The energy for this process is provided by ATP molecules. Thus both sodium and potassium ions are essential for living organisms.
Lithium Symbol – Li Atomic no. - 3 Atomic Weight – 6.94u Electronic Configuration - 1s22s1 Group no. – 1 Period no. – 2 Group name – Alkali Metals Block name – ‘s’ Standard State(298 K)- Solid Color – Silvery-white/grey Classification - Metallic
Anomalous Properties High melting & boiling point. Much harder than other alkali metals. Reacts with oxygen least readily to form normal oxide(E.g. Li2O), whereas other alkali metals form peroxides and superoxides(E.g. MO2,M2O2). 4Li + O2→ 2Li2O Unlike other alkali metals lithium reacts directly with carbon to form an ionic carbide.
The carbonates, hydroxides and nitrates of lithium decompose on heating unlike those of other alkali metals which are somewhat stable towards heat. 4LiNO3→ 2Li2O + 4NO2 + O2 2LiCO3→ 2Li2O + CO2 2LiOH → Li2O + H2O LiOH is a weaker base than hydroxides of other alkali metals. Unlike elements of group 1, Lithium forms nitride with nitrogen. 3Li + N → Li3N Lithium halides are have more covalent nature than halides of other members of group 1.
Due to appreciable covalent nature, the halides and alkyls of lithium are soluble in organic solvents. • Li+ has very high hydration energy and charge/radius ratio, therefore it acts as an excellent reducing agent in solution. Li + Li + Li + Li +
Small size of atom results in relatively high cohesive properties associated with relatively strong inter-metallic bonding; large atoms usually form weak bonds.
Diagonal Relationship B C The properties of lithium are quite different from the properties of other alkali metals. On the other hand, it shows greater resemblance with magnesium, which is diagonally opposite element of it. Similarly properties of Beryllium & Boron represent that of Aluminium & Silicon respectively. The main reasons for the anomalous behavior of lithium are -:
The Reasons -: (i) The extremely small size of Lithium & its ion. (ii) Greater polarizing power of lithium ion ( Li+), due to its small size which result in the covalent character in its compounds. (iii) Least electropositive character and highest ionization energy as compared to other alkali metals. (iv) Non availability of vacant d-orbitals in the valence shell.
Some More ExamplesExamples For Diagonal Relationship Li and Mg form only normal oxides whereas Na forms peroxide and metals below Na, in addition, forms superoxide. Li is the only Group 1 element which forms nitride, (Li3N). Mg, as well as other Group 2 elements, also form nitride. Lithium carbonate, phosphate and fluoride are sparingly soluble in water. The corresponding Group 2 salts are insoluble. (Think lattice and solvation energies). Both Li and Mg form covalent organometallic compounds. LiMe and MgMe2 (of Grignard reagents) are both valuable synthetic reagents. The other Group 1 and Group 2 analogues are ionic and extremely reactive (and hence difficult to manipulate).
THE ALKALI METALS ARE HIGHLY REACTIVE. CAUSING CONTRIBUTING FACTORS ARE LARGE SIZE AND LOW IONIZATION ENTHALPY. THE REACTIVITY OF THESE METALS INCREASES DOWN THE GROUP.
CHEMICAL PROPERTIEES 1. REACTIVITY TOWARDS AIR -
ALKALI METALS TARNISH IN DRY AIR DUE TO FORMATION OF THEIR OXIDES. • 2. THEY BURN IN OXYGEN VIGOURSLY. • 3. THEY REACT WITH MOISTURE FORMING HYDROXIDES.
EX- 4LI+O2 2LI2O 2NA+O2 NA2 O2 M+O2 MO2 OXIDATION STATE- +1
LITHUIM IS AN EXEPTION REACTING DIRECTLY WITH NITROGEN OF AIR TO FORM THE NITRIDE. DUE TO THEIR HIGH REACTIVITY TO AIR AND WATER THEY ARE KEPT IN KEROSENE OIL.