s-Block Chemistry Flashcards
Ionic model
Ions re assumed to be hard spheres with fixed sizes held together by electrostatic interactions, the radii of spheres are known as ionic radii.
Internuclear distance
Can be measured by x-ray crystallography but it is hard to determine individual radii as hard to know when electron density stops and starts. Ionic radii change depending on the compound.
Trends in ionic radii
Anions>cations
Increase down a group; Increase in PQN
Dicationmonoanion
Lattice Enthalpy
Enthalpy change for the conversion of 1 mol of the ionic solid into the gaseous state
ΔU in the ionic model:
-ΔlattH° in the ionic model comes from electrostatic interactions between ions, by considering these interactions in an ionic solid you can obtain a theoretical value
ΔU = -(z- z+ e^2)/(4 π εo r)
Madelung constant
in a crystal there are repulsive and attractive interactions in all directions to account for this the Madelung constant is introduced as well as Avogadro’s number
ΔUlatt = +(A Na z- z+ e^2)/(4 π εo r)
Lattice energy is the -ve of this potential so the +ve sign is now included
Born-Lande Equation
The prior expression assumes ions are point charges; additional short range forces need to be included from overlapping electron clouds, these increase as)decreases.
+(A Na z- z+ e^2)/(4 π εo r)(1-1/n)
The born exponent is an average of an n value obtained by comparing the compressibility of a solid based on its occupied orbitals.
ΔU/ΔH
All prior calculations give internal energy changes, ΔU, however ΔH is heat change at constant pressure, they are related by: ΔH = ΔU + pV; the difference is so small hence can be ignored
Calculating the Madelung constant
The first attractive energy equation is used, times by the # of nearest neighbours, r is calculate during extended Pythagoras theorem. Include sign +/- !!!
d = (2r^2)^1/2 d = (3r^2)^1/2
ΔU = ΔU1 + ΔU2 + ΔU3 …
The terms alternate in sign and are of high but decreasing magnitude.
Deviations from the ionic model
+ve charges distort electron clouds of anions, large and small charged ions are more easily polarisable. highly charged small cations bonded to large charge diffuse anions have the greatest degree of covalent character.
The BL eq. underestimates ΔUlatt high covalent character compounds.
ΔUlatt(BH) ≈ ΔUlatt(BL) -> low covalency
ΔUlatt(BH) > ΔUlatt(BL) -> high covalency
Kapustinskii Equation
A simplification of the Born-Lande equation;
ΔUlatt = (k ν z+ z-)/(r+ + r-)
Where k = 107900, ν = no. ions in the formula units
Thermochemical radii of polyatomic ions can be calculated in this way
van Arkel-Ketelaar triangles
Predict the bonding type for a binary compound;
Metallic ^^ Ionic ^^ Covalent
X-axis - Average electronegativity
Y-axis - Difference in electronegativity
When both χ’s are high –> covalent, when high Δχ –> ionic
Hydrogen
H2 is colourless and odourless, MP: -259°C, BP: -253°C (low/weak IM forces), low Mr, lowest density of all gases (0.082gdm-3). Unreactive at RTP (H-H 436JKmol-1) with the exception of O2, F2 and Cl2.
Production of H2
Steam reforming: methane + water -> carbon monoxide + 3hydrogen (NiO/850°C)
Shift reaction: CO + H2O -> CO2 + H2 (Fe/450°C)
Electrolysis of H2O: 2H20 -> 2H2 + O2
Use of hydrogen
Mainly used in the Haber process (53%), in the petrochemical industry, for extraction of metals from ores and in ethanol production (17%)
Hydrides
Hydrogen forms binary compounds called hydrides with most elements; χH>χR (R is electropositive) the H atom is hydride and has an oxidation state of -1. If χH
Hydridic Hydrides
Hydridic hydrides are formed by group 1 & 2 and are ionic structures, strongly basic
RH + H2O -> ROH + H2
Reactivity with water increases down the group hence they are kept under inert atmospheres.
Protic hydrides
Covalent not ionic as ionisation (1312KJmol-1) »_space; electron gain (-73KJmol-1) hence H+ difficult to produce. Only formed when dissolved in a solvent that can solvate H+; the compound here acts as an acid
Non-polar hydrides
(χH≈χR) may have a small dipole: i.e. in CH4 H=δδ+ and in B2H6 H=δδ-
Nomenclature of hydrides
Hydridic hydrides -> hydridides
Protic hydrides/Non-polar hydrides -> R-ane
Group 16&17 ->hydrogen R-ide
Trivial names - Ammonia, water
Covalent Hydrides
Electron precise compounds - All valance electrons are involved in bonding (group 14)
Electron deficient compounds - 3-centre-2-electron bonds (group 13)
Electron rich compounds - Not all electrons on the central atom are involved in bonding (lone pairs) hence can act as lewis bases (groups 15-17)
Acidity of hydrides
Increases across a period (Δχ increases) and increases down group (decrease in BDE, decreasing attraction between X and H3O+)
HX + H2O –> H3O+ X-
BDE of hydrides
Increases from left to right due to an increase in ionic contribution as χX increases. Decrease down group as valance orbitals get larger and more diffuse as PQN increases therefore ns np interactions with H 1s are reduced.
Isotope effects of hydrides
Protium/deuterium/tritium
H^2 Has twice the mass of H^1 hence its properties will vary considerably, D2O ice sinks as it is more dense,
BDE of D2 443>H2 436 therefore more energy is needed to break X-D that X-H; the lowest vibrational energy state for D2 is lower than for H2
Group 1 general properties
Li, Na, K, Rb, Cs, Fr (radioactive longest lived isotope 22mins, minor component in uranium minerals
Low MPs, BPs, ΔaH° -> weak metallic bonding
Low densities -> large atomic radii, open body-centered cubic structures
Preparation of group 1 metals
Li and Na electrolysis of molten chlorides in a downs cell, CaCl2 added to lower the MP
2Na+ + 2e- -> 2Na (reduction)
2Cl- -> Cl2 + 2e-
K, Rb, Cs reduction of molten salts with Na at high temps
KCl + Na -> K NaCl
eq lies to the left therefore K (volatile) removed by fractional distillation
Group 1 oxides
All burn in air to form oxides: Li (O2- oxide), Na (O2^2- peroxide), K (O2- superoxide)
All are basic and fact with water to form hydroxides, H2O2 and oxygen as you descend the group
Large anions are stabilised by large cations; G1 cations increase down the group therefore are better at stabilising large peroxide and superoxide anions
Li2O2 decomposes on heating whereas Na2O2 is stable to it (Born hater cycle using ΔG = ΔH-TΔS)
Group 1 suboxides
Rb and Cs burn in limited amounts of O2 to form coloured suboxides i.e. Rb9O2; delocalised electrons cause metallic behaviour
Group 1 hydroxides
Metal reacts with water to form hydroxide and hydrogen gas. ΔH is -ve, reactivity increases down the group. Li,Na, K less dense than H2O so react on the surface, the metal melts and, for K the H2 gas ignites. Rb and Cs are more dense so sink and rect explosively. NaOH produced industrially by electrolysis of NaCl (chloroalkali process)
2NaCl + 2H2O -> 2NaOH + H2 + Cl2
Group 1 halides
Colourless ionic solids, high MPs, rock salt structure
Group 1 ethynides
All metals react to with ethyne in liquid ammonia to form the NH2^- mono cation or the C2^2- dianion.
Group 1 ethynides decompose in H2O to form LiOH and C2H2
Group 1 nitrides
Li only metal to form a stable binary nitride (stability decreases down the group) Li3N decomposes in water to form LiOH and NH3
Oxoanions
Group 1 metals form salts i.e. NO3-, CO3^2-, SO4^2-. Group 1 nitrates (NO3-) decompose on heating to nitrites (NO2-) and oxygen. For lithium decomposition gives 2Li2O + 4NO + 3O2
Nitrates become more stable down the group due to the decreasing difference between the lattice enthalpies of nitrate and nitrite making decomposition less favourable..
Compounds with small anions become less stable down the group (decrease in attic enthalpies of reactants). Compounds with large anions increase in stability down the group (decrease in lattice enthalpies of decomposition products).
Solubilities of group 1 metals
A compound is soluble if ΔsolG < 0 as entropy cannot be ignored
ΔsolG = ΔlattG + 2ΔhydH
Large anions -> solubility decreases
Small anions -> solubility increases
Small cation and small anion/large cation and anion -> INSOLUABLE ΔlattG +ve and dominates as everything either all high or all low respectively
For large and small combination ΔhydG of the smaller ion dominates hence -ve and dominates.
Group 1 coordination chemistry
M+ are lewis acids so can react with lewis base ligands. Group 1 have low charge density and ar relatively large hence are weak lewis acids. Coordinate weakly to H2O, decreases down the group.
Crown ether coordination; relationship between cavity size, canonic radius and stability -> optimum spatial fit.
[n1]crown-n2 where n1=atoms in ring and n2=no oxygen atoms
Cryptands are also size selective and have a central cavity
