Ionic Bonding and s-Block Chemistry- Group 1 Flashcards
Describe the state of the first 5 members of group 1
- Metallic solids at room temperature and pressure
2. Melting point of caesium is low enough for it to melt in hot weather
Describe Francium
- Radioactive
- not been isolated as the pure element
- Found as a minor component in uranium minerals
- The longest-lived isotope has a half life of only 22 minutes
What is the chemistry of group 1 elements dominated by
- +1 Oxidation state
Describe the general physical properties of group 1
- Low melting/boiling points and enthalpies of atomisation reflect relatively weak metallic bonding
- Low densities due to large atomic radii and relatively open body-centred cubic structures
How is lithium and sodium prepared
- Industrially by electrolysis of their molten chlorides
2. For sodium, this is carried out in a Downs cell, and CaCl2 is added to reduce melting point
Show the equations that take place in the electrolysis Na
- Reduction at cathode- 2Na+ + 2e- –> 2Na (l)
2. Oxidation at the anode- 2Cl- –> Cl2 (g) + 2e-
How are potassium, rubidium and caesium metals prepared
- By reduction of their molten salts with sodium at high temperatures
Show equation for preparation of K from KCl and say the driving force
- KCl + Na K + NaCl
- Potassium is more reducing than sodium so the equilibrium lies to the left hand side
- More volatile potassium is obtained by fractional distillation, which displaces the equilibrium to the right-hand side
Describe what happens to group 1 metals when they burn in air
- Burn to form oxides
- The major product depends on the metal
- But can all form an oxide, peroxide and superoxides under appropriate conditions
What does lithium burn to form and show equation
- Lithium oxide- Li2O
- 4Li (s) + O2 (g) –> 2Li2O (s)
- O 2-
What does sodium burn to form and show equation
- Sodium peroxide- Na2O2
- 2Na(s) + O2 (g) –> Na2O2 (s)
- O2 2-
What does Potassium and heavier metals burn to form and show equation
- Form superoxides (MO2)
- K(s) + O2 (g) –> KO2 (s)
- O2 -
How do group 1 oxides react with water
- All basic
2. React with water to give hydroxides
Show equations of Li2O, Na2O2, KO2 reactions with water
- Li2O (s) + H2O (l) –> 2LiOH (aq)
- Na2O2 (s) + 2H2O (l) –> 2NaOH (aq) + H2O2 (aq)
- 4KO2 (s) + 2H2O (l) –> 4KOH (aq) + 3O2 (aq)
What can KO2 reaction be used for
- Used in submarine breathing systems to generate O2
2. The KOH product absorbs CO2
Why does the type of oxide produced by combustion change going down the group
- Large anions are stabilised by large cations
- The ionic radius of group 1 cations increase down the group
- And the larger cations are better at stabilising the large peroxide and superoxide ions with respect to decompositions into oxide and oxygen gas
- Lithium peroxide decomposes on heating but sodium peroxide is stable to heating
What else can caesium and rubidium burn to form
- In limited oxygen can form intensely coloured compounds called suboxides
Describe suboxides e.g. Rb9O2
- The metal oxidation state in a suboxide appears to be less than +1- Not the case
- The additional electrons are actually delocalised across the whole structure
- So the formula for Rb9O2 could be written (Rb+)9(O2-)2(e-)5
- These delocalised electrons give rise to metallic behaviour
What is produced when group 1 metals react with metals
- Give hydroxide and hydrogen gas
Describe th reaction of group 1 metals with water
- Very exothermic
- Violence of reaction increases going down the group
- Li, Na, K are all less dense than water so they react on the surface
- All the reactions except Li, the reactions are exothermic enough to melt the metals
- Reaction with K is vigorous to ignite the hydrogen product
- Rb and Cs are denser than water so they sink beneath the surface and react explosively
How is NaOH prepared industrially
- By electrolysis of NaCl solution known as the chloralkali process
- 2NaCl (aq) + 2H2O (l) –> 2NaOH (aq) + H2 (g) + Cl2 (g)
Describe group 1 halides
- Colourless ionic solids with high melting points
How is NaCl obtained
- Mining naturally occuring deposits
2. Evaporation of sea water
Describe group 1 ethynides
- Group 1 metals all react with ethyne in liquid ammonia to form ethynides
- Contain HC2- monoanion or the C22- dianion
Show the equations for the formation of Lithium ethynides
- 2Li(s) + 2HC=-CH (g) –> 2Li+C=-CH- (s) + H2 (g)
- 2Li(s) + HC=-CH (g) –> (Li+)2C=-C 2- (s) + H2 (g)
=- Is a triple bond
How do group 1 ethynides decompose e.g. Li2C2
- Decompose in water to form hydroxide and ethyne
2. Li2C2 (s) + 2H2O (l) –> 2LiOH (aq) + C2H2 (g)
Which is the only group 1 metal to form a stable binary nitride
- Lithium
How is Lithium nitride prepared
- It is prepared from the reaction of lithium with nitrogen at high temperature and pressure
- 6Li (l) + N2 (g) –> 2Li3N(s)
How is lithium nitride decomposed
- Decomposes in water to form lithium hydroxide and ammonia
2. Li3N (s) + 3H2O (l) –> 3LiOH (aq) + NH3 (g)
What interest is there in Li3N
- Lithium Nitride shows high Li+ ion conductivity
2. Is being investigated for use in Li- ion batteries
Why does Li3N form but not sodium nitride
- DfH for lithium nitride is negative so its formation from the elements is favourable
- The term in the Born-Haber cycle with the largest magnitude is the lattice enthalpy which is very high due to high charge on N3- and small size Li+
- DlattU are different between Li and Na as Na has a higher radius so is less positive so -DLattU is a lot less exothermic
What else do group 1 metals form salts with
- With oxoanions such as nitrates, carbonates and sulfates
How do group 1 nitrates decompose
- MNO3 decompose on heating to the nitrites MNO2
- 2MNO3 (s) –> 2MNO2 (s) + O2 (g)
- But with lithium decomposition gives the oxide
- 4LiNO3 (s) –> 2Li2O (s) + 4NO (g) + 3O2 (g)
How does the stability of nitrates change down group 1
- Nitrates become more stable with respect to decomposition as the group is descended
- The increasing stability is due to decrease in difference between lattice enthalpies of the nitrate and nitrate which makes decomposition less favoured
What is the general rule for stability of compounds with small anions
- Less stable down the group
- This is due to decrease in lattice enthalpies of these compounds a the group is descended
- When these compounds decompose on heating, they decompose to the elements
What is the general rule for stability of compounds with large anions
- More stable down the group
2. Due to decrease in lattice enthalpies of the decomposition products
Describe the solubility of group 1 salts
- Soluble in water
- Salts containing large anions, the solubility generally decreases down the group
- Salts with small anions, the solubility increases down the group
State some large anions
- Chlorides, Bromides, Iodides and Nitrates
State some small anions
- Fluorides and hydroxides
What does the solubility of a compound depend upon
- The relative magnitudes of the lattice Gibbs energy and the Gibbs energy changes of hydration of the ions
- Gibbs energies are used instead of enthalpies as entropies can’t be ignored
What is the equation for the solubility of a compound and what do different values mean
- DsolG(MX) = DlattG(MX) + DhydG(M+) + DhydG(X-)
- A compound is soluble if DsolG is negative
- Both DlattG and DhydG are large but with opposite signs.
- This means DsolG is small which is why entropies cannot be ignored
Describe the solubility if there is a small cation and small anion
- DlattG(MX) -very high
- -DhydG(M+) - high
- -DhydG(X-)- high
- DlattG(MX) dominates so DsolG(MX) is positive and the compound is insoluble
Describe the solubility if there is a small cation and large anion
- DlattG(MX) - high
- -DhydG(M+) - very high
- -DhydG(X-)- high
- DhydG(M+) dominates so DsolG(MX) is negative and the compound is soluble
Describe the solubility if there is a large cation and small anion
- DlattG(MX) - high
- -DhydG(M+) - low
- -DhydG(X-)- very high
- DhydG(X-) dominates so DsolG(MX) is negative and the compound is soluble
Describe the solubility if there is a small cation and large anion
- DlattG(MX) - high
- -DhydG(M+) - very high
- -DhydG(X-)- high
- DlattG(MX) dominates so DsolG(MX) is positive and the compound is insoluble
What is the basis of coordination chemistry
- Metal ions are Lewis acids, so they are able to interact with Lewis bases
- This type of interaction is the basis of coordination chemistry, in which the Lewis bases are called ligands
How strong is the coordination with group 1 ions
- They are singly charged and relatively large so have a low charge density
- As a result, they are weak Lewis acids and, unlike many other metals, they coordinate weakly to simple ligands such as water
- The coordination becomes weaker down the group as the charge density on the cation decreases
What is a crown ether
- First prepared accidently
- Cyclic polyethers
- There is a relationship between the cavity size, cationic radius and the stability of the resulting complex- optimum spatial fit
What is valinomycin
- Naturally occurring macrocylic antibiotic
- Neutral molecule made up from 6 amino acids and 6 hydroxy acids alternating in a ring
- It is very selective for K+ as Na+ is too small t interact with 6 oxygen atoms
- The hydrophobic isopropyl groups allow valinomycin to transport K+ ion across cell membranes and this damages bacteria
What are cryptands
- Another class of ligands that bind to group 1 cations
- They form more stable complexes than crown ethers as less rearrangement of the ligand is needed to coordinate to the cation
- Like crown ethers, cryptands are size-selective and they form most stable complexes with cations that fit best into the central cavity
What happens when group 1 metals react with liquid ammonia
- Dissolve to give dark blue solutions
- Unlike the reaction of sodium with water, the solvent is not being reduced
- Instead the metal atoms ionise the cations and electrons, both of which are solvated by liquid ammonia
Give equation for solvation of Na
- Na (s) –> Na+ (solv) + e- (solv)
What produces the colour of the solution when group 1 metals are reacted with liquid ammonia
- Dark blue colour is due to the solvated electrons
2. These electrons sit in cavities formed by groups or NH3 molecules
What do the solvated electrons act as
- Free electrons
2. So the solutions are strong reducing agents
Show how C60 is can be reduced by group 1 metals
- C60 is reduced to [C60]- anion on reaction with rubidium in liquid ammonia
- C60 + 3Rb+ (solv) + 3e- (solv) –> Rb3[C60]
Show how Nickel(II) complex [Ni(CN)4]2- is reduced to a nickel (0) complex
- Reduced by sodium in liquid ammonia
2. [Ni(CN)4]2- + 2Na+ (solv) + 2e- (solv) –> [Ni(CN)4]4-
What happens when there is a high concentration of metal reacting with liquid ammonia
- The blue colour turns into a metallic bronze and the liquid ammonia solution begins to conduct electricity
- Electrons are delocalised throughout the solution in a similar manner to the delocalisation of electrons in a solid metal
What happens to solvated electron solutions e.g. with Na
- THey are thermodynamically unstable
- After several days ammonia is reduced to form sodium amide and hydrogen
- 2Na(s) + 2NH3 (l) –> 2NaNH2 (solv) + H2 (g)
