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6.0 Main Group Elements: Typical Chemistry and Compounds

6.1 Hydrogen: READING: Chapter 9, Sections 9.2, 9.4, 9.5, 9.9 and Chapter 4, Section 4.2

• 99.9% of the known universe is empty. Of the ~0.1%

that is not, ~91% is hydrogen in the form of interstellar

clouds and stars.

• H (and He) are unique in having only the 1s orbital

available for chemistry.

• H is the only element known to form compounds with

all other elements!

• Under STP conditions,

dihydrogen (H2) is a gas.

However, under high pressure,

hydrogen can form a metal

similar to sodium. The cores of

Jupiter and of Saturn are believed

to contain metallic hydrogen.

• E-H compounds are called hydrides. When E is a metal, ionic (M+H¯) salts or conductive

metal solids are formed. When E is a p-block element, covalent hydrides are formed.

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• H¯ (hydride ion) is a strong Lewis base.

• Hydrogen’s cation is a proton (H+), which is a strong Lewis acid.

Let’s look at the Carbon Group hydrides (C, Si, Ge, Sn, Pb)

R3C-H The C-H bond tends to be strong and essentially non-polar.

R3Siδ+-Hδ- Silanes are very strong reducing agents (sources of H-)!

Only SiH4 and Si2H6 are indefinitely stable at 25 °C.

Silanes are spontaneously flammable in air.

e.g., Si4H10 + 13/2 O2 → 4 SiO2 + 5 H2O

Silanes are stable in water and dilute mineral acids, but rapidly hydrolyze

in base. e.g., Si2H6 + (4+2n) H2O → 2 SiO2•nH2O + 7 H2

Silanes react explosively with halogens at 25 °C.

R3Ge-H The Ge-H bond tends to be moderately strong and essentially non-polar.

Germanes (GeH4, Ge2H6, Ge3H8, etc.) are less flammable than silanes and

resistant to hydrolysis, even at 30% NaOH.

Q. Why are germanes closer in reactivity to carbanes than to silanes?

R3Snδ+-Hδ- Stannane (SnH4) is easily oxidized; a useful reducing agent.

R3Pbδ+-Hδ- The existence of plumbane (PbH4) is dubious.

• 1H NMR is particularly useful for studying E-H compounds.

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• Hydrogen bonding is one of the key concepts to understanding chemical structures, solid

state structures and physical properties of many compounds.

• Trends in boiling points of molecular hydrides can be correlated, in part, to hydrogen

bonding. Note also the correlation to electronegativity!

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Protic Acids and H-bonding

The Arrhenius/Ostwald Concept (1890)

• “An acid is a substance that tastes sour and dissociates protons in aqueous solution.”

The Brønstedt-Lowry Concept

• “Proton donors are acids. Proton acceptors are bases.”

• Gives rise to the concept of conjugate acids and conjugate bases.

H:A + :B → :A¯ + H:B+

acid base conj. base conj. acid

• The Brønsted-Lowry concept also underlies

hydrogen bonding:

• Let’s take a closer look at FHF¯ from an MO

perspective (again) as a typical hydrogen bond.

Brønstedt acid: LUMO = 1s orbital of H+

Brønstedt base: HOMO = most often like a

lone pair (large lobe)

• 3-center-4-electron bond

• hydrogen bonds are typically asymmetric

Source: Purcell +

Kotz, Inorganic

Chemistry, Holt-

Saunders, 1971

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• Overlap of the donor orbital D with σ*AH

1) stabilizes σnAHD and 2) destabilizes σ*

AHD → leads to overall bonding interaction • Overlap of the donor orbital D with σAH

1) stabilizes σAHD and 2) destabilizes σnAHD

→ no overall bonding interaction • The interaction of the σ*

AH orbital is crucial!

The more electronegative A (polarized bond towards A), the stronger the overlap between σ*

AH and D. This is why only O-H, N-H, and F-H are expected to form strong hydrogen bonds!

What makes an acid acidic and a base basic?

• Why is HCl(aq) an acid and NH3(aq) a base? Let’s have a close look:

HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq)

NH3(g) + H2O(l) → NH4+(aq) + OH-(aq)

• Breakdown of these reactions into elementary steps (Born-Haber-Cycles):

Acid reaction:

ΔH = -BEHCl + IP(H) + EA(Cl) + ΔHSolv(H+) + ΔHSolv(Cl-)

= 102.2 + 313.5 - 83.3 -270.3 -77.2 = -15.1 kcal/mol

• The reaction of HCl with H2O is exothermic and driven largely by ΔHSolv(H+).

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Base reaction:

ΔH = ΔHvap - BEOH + IP(H) + EA(OH) + PA(NH3) + ΔHSolv(NH4+) + ΔHSolv(OH-)

= 10 + 119 + 313 - 42 - 207 - 90 -110 = -7 kcal/mol

where PA = proton affinity as defined on the next page …

• The reaction of NH3 with H2O is exothermic and driven by the solvation energies.

• The above procedure is of general applicability:

Need to know various thermodynamic parameters to do the calculations, which can be

independently measured.

CONCLUSION: Adicity/Basicity is f(solvent)

HOMEWORK: Write down the expression for ΔH for the same reaction in liquid NH3?

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Proton Affinity (PA) and Hydrogen Affinity (HA): Heterolytic vs. Homolytic Cleavage

Proton Affinity:

Def.: XH+(g) ← X(g) + H+(g) Heterolytic (both electrons in bond go to X)

(analogous to electron affinity) NOTE: gas phase!

• Experimentally determined by ion-cyclotron resonance. (high res. mass spec.)

Hydrogen Affinity:

Def.: XH+(g) ← X(g)+ + H•(g) Homolytic (one electron each to X and H)

Relationship between PA and HA:

PA = HA + (IPx - IPH), i.e. PA and HA are related by the difference in IP between X and H,

where IPH = 313.6 kcal/mol - always the same for all Brønstedt acids.

• Some real life data and examples:

Source: W.W. Porterfield, Inorganic Chemistry, Addison-Wesley, 1984.

• The PA of any species is first a function of charge:

Lowest PA of anion in table: I- = 313 kcal/mol

Highest PA of neutral species in table: NH3 = 207 kcal/mol

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Consequences: HCl(g) + H2O(g) → H3O+(g) + Cl-(g) ΔH = 333-164 kcal/mol … does actually not happen in the gas phase! i.e. equilibrium is far on the left side. As shown before the reaction is driven by ΔHSolv(H+) • Some notable trends that do correlate with solution behaviour:

Acidity HI > HBr > HCl > HF Acidity H2Te H2Se H2S H2O

pK1 2.63 3.89 6.99 15.7 pK2 ~11 ~11 12.9 ~29

Can you explain these trends using some of the data in the table? • Illustrating the difference between HA and PA …

Data Source: D.H. Aue et al., JACS, 97, 4136, (1975), ibid., 98, 318, (1976). sp3 nitrogen stronger donor towards H+ (more p character!) but… sp2 nitrogen forms stronger bond with H• (higher s-character of hao → stronger bond). PApy - PApip = (HApy - HApip) - (IPpy - IPpip) = -5 = 19 - 24 where ΔHA reflects the difference in HOMO/LUMO overlap (better overlap with sp2) and ΔPA reflects the actual Brønstedt basicity, which is related to the HOMO/LUMO energy match between the base (HOMO) and the proton (LUMO). → Piperidine is the stronger base.

HOMEWORK: Exercises 9.1, 9.3, 9.5 – 9.11, 9.13, 9.14, 9.16, 9.19

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Solvent Leveling

• In any solvent that can autodissociate, the cation is acidic and the anion is basic:

2 H2O H3O+ + OH- K ≈ 10-14

2 NH3 NH4+ + NH2

- K ≈ 10-27

2 H2SO4 H3SO4+ + HSO4- K ≈ 10-4

3 HF HFH+ + FHF- K ≈ 10-12

2 SO2 SO2+ + SO32- inferred, no direct evidence

NOTE: This concept is not restricted to protic solvents (e.g., SO2 )

In any given solvent: • solutes that increase the cation concentration are acids • those that increase the anion concentration are bases. The strength of an acid or base is function of the solvent it is in. EXAMPLE: No acid stronger than H3O+ can exist in water.

No base stronger than OH¯ can exist in water. H2O levels all acids to a maximum strength equal to that of H3O+.

HX(sol) + HSolv(l) H2Solv+(sol) + X-(sol)

All acids HX with pKa < 0 have the acidity of H2Solv+ in the solvent HSolv.

H2O levels all bases to a maximum strength equal to that of OH¯.

:B(sol) + HSolv(l) HB+(sol) + Solv(l)¯

All bases with pKb < 0 have the basicity of Solv¯ in the solvent HSolv.

• pKa + pKb = pKsolv Always!

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• Any acid is leveled in HSolv if pKa < 0.

• Any base is leveled in HSolv is pKa > pKsolv

• For any solvent HSolv there is a “window” from pKa = 0 to pKa = pKsolv in which acids or bases are not leveled:

Source: Shriver & Atkins, Inorganic Chemisty, 3rd Edition, Freeman. HOMEWORK: Exercises 4.2 – 4.4, 4.6

6.2 Group 1 (Alkali Metals): Li, Na, K, Rb, Cs, Fr*

READING: Chapter 10, Intro and Sections 10.3 – 10.12

• Monovalent metals with electron configuration ns1.

Very electropositive and reducing.

• React violently with acids (e.g., with H2O to form

alkaline hydroxides MOH, e.g. NaOH, KOH.)

• Reactivity increases down the group with inc. atomic radius (decreasing I.E.)

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NIST-F1, the American primary time and frequency standard, is a

cesium fountain atomic clock developed at the NIST laboratories in

Boulder, Colorado.

Typical compounds:

• Salts: NaCl (cooking salt), KBr (IR medium), CsF (fluorination agent in organic chemistry)

• Hydrides such as LiH (lithium hydride; drying agent), NaH (useful non-nucleophilic base

and reductant: e.g., NaH(s) + NH3(l) → NaNH2(am) + H2(g) • Hydroxides such as NaOH, KOH, as mentioned above • Azides such as NaN3 (useful reagent for organic and inorganic chemistry) • Organometallic compounds, primarily of Li.

Lithiating agents are commonly used in organic synthesis. Organolithium reagents are often pyrophoric (ignite spontaneously in air below ca. 55 °C).

nBuLi exists as clusters in the solid state. tBuLi is significantly more reactive!

e.g. tetramer of n-BuLi Fume hood after a fatal tBuLi fire at UCLA in December 2008.

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Solutions in Liquid Ammonia Q. What happens when you put sodium in water? Q. Why doesn’t the same thing happen in liquid ammonia?

NH3(l) + Na(s) → Na+(am) + NH2

‾ + ½ H2(g)

A. This reaction is kinetically inhibited…but can be catalyzed (causing a rather violent explosion) by trace amounts of transition metals (e.g. Fe2+/3+)

Therefore this is possible:

NH3(l) + Na → Na+(solv) + e-(solv) + NH3(l)

• The blue to bronze solution of electrons in ammonia is

kinetically, but not thermodynamically, stable. • What does thermodynamically unstable and kinetically

inhibited mean? Thermodynamic parameters: (standard state @ 25 oC 1 M NH4

+ or NH2- solutions)

The relevant half-cell reactions H2/NH3 (reduction of NH3) and NH3/N2 (oxidation of NH3) in liquid ammonia are (by convention always written as reductions): Acidic medium (1 M NH4

+): NH4

+ + e- → NH3 + 1/2 H2 Eo = 0 V (Reduction, NH4+ cation is reduced)

3 NH4+ + 1/2 N2 + 3 e- → 4 NH3 Eo = 0.04 V (Oxidation, neutral NH3 is oxidized)

Basic medium (1 M NH2-):

NH3 + e- → NH2

- + 1/2 H2 Eo = - 1.59 V (Reduction, neutral NH3 is reduced) 2 NH3 + 1/2 N2 + 3 e- → 3 NH2

- Eo = - 1.55 V (Oxidation, anionic NH2- is oxidized)

These redox potentials mean that:

in acidic solution, any solute that is a stronger oxidant than N2 (E > 0.04) should oxidize NH3 to N2 and any solute with E < 0.0 V should reduce NH4

+ to NH3 and H2. in basic solution, any solute with E > - 1.55 V should oxidize NH2

- to NH3 and any solute with E < - 1.59 V should reduce N2 to NH2

-.

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• In both cases there is only a small 0.04 V window of redox stability for the acidic or basic NH3 solution, i.e. only very few solutes would not oxidize or reduce NH3.

But the reduction potential of Na in NH3 is -1.85 V!

So how is it still possible to dissolve Na in NH3(l)???

Theoretically (solid lines in diagram): Redox stability

only in narrow 0.04 V window.

Practically (dashed lines in diagram): Overpotentials

result in redox stability from about + 1.00 to - 1.00 V.

• Oxidizing (MnO4-, O3, etc.) and reducing (Na, K,

etc.) compounds can be dissolved in NH3(l)

The overall solubility of alkali metals in NH3(l) at -35 oC are quite substantial: Metal Solubility (g/kg) in NH3(l) Li 15.7 Na 10.9 K 11.9

Source: Purcell&Kotz, Inorganic Chemistry, Holt-Saunders, 1977. HOMEWORK: Exercises 10.1, 10.3 - 10.6 6.3 Group 2 Alkaline Earth Metals: Be, Mg, Ca, Sr, Ba, Ra* READING: Chapter 11, Intro and Sections 11.2 – 11.10; Chapter 4, Section 4.10 & 4.11 • Divalent; Electropositive; Smaller atomic radii than Group 1 of same period, which accounts

for higher I.E.s. Typical compounds: • Salts: CaCl2 (drying agent), CaF2 (fluorspar), MgSO4 (drying agent), CaSO4 • xH2O

(gypsum, alabaster), BaSO4 (contrast agent for medical X-ray), CaCO3 (Rocky Mountains), Ca5(PO4)3F (apatite), Ca5(PO4)3OH (hydroxy apatite, a key component of vertebrate bones).

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• Readily forms oxide MO, which hydrolyze to give dibasic hydroxides M(OH)2, e.g., Ca(OH)2 (lime); Mg(OH)2 (“Milk of Magnesia”).

• Carbides MC2: CaC2(s) is actually an

acetylide (i.e., calcium salt of the C≡C2- anion) and is still used by cavers as a light source (reaction with moist air causes a slow production of acetylene which can be burned to give a long-lasting flame).

• Organometallic compounds:

Grignard reagents RMgCl and RMgBr

Some notes on specific elements: • Beryllium (Be) and its compounds are highly toxic.

Used for high-grade astronomical mirror coatings, and for various nuclear applications such as subcritical nuclear ignition devices and as a neutron booster in “clean” neutron bombs.

• Radium (Ra) was one of the first radioactive elements identified by Antoine Henri Becquerel,

Marie and Pierre Curie (Nobel Prizes 1903 and 1911 – see http://www.nobel.se).

Marie Curie Pierre Curie Henri Bequerel

Standard units of radioactive decay: 1 Bequerel = 1 decay per second; 1 Curie = 3.7 × 1010 Bq

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Hard-Soft Acid-Base Theory (HSAB): Predicting stabilities of adducts H+ is the key electron pair acceptor (acid) in the Brønsted acidity scheme. Lewis acidity and basicity: Any species that accepts an electron pair is an acid. Any species that donates an electron pair is a base. Consequence: Monocations of alkali metals and dications of alkaline earth metals are Lewis acids! • The concept of hard and soft acids and bases is of universal applicability in chemistry

(inorganic, organic, bio, physical, toxicology) and was developed by R.G. Pearson.

• Generalized acid-base reaction (Lewis concept):

• How does the nature of A/A' and B/B' influence the equilibria?

How can we make some predictions?

• Empirically, based on experimentally determined adduct stabilities, acids and bases can be divided into to two classes:

Class A Class B Hard Soft

high charge density - not very polarizable low charge density - polarizable high electronegativity (base) low electronegativity (base)

ΔEHOMO/LUMO large ΔEHOMO/LUMO small acid/base tends to bind to other another

hard base/acid acid/base tends to bind to other another soft

base/acid predominantly electrostatic predominantly covalent

very polar A-B bond less polar A-B bond no backbonding synergistic back-bonding possible

valence electrons not easily excited valence electrons easily excited typically not very toxic typically very toxic

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• Empirical classification of bases:

Class A: Hard Bases Class B: Soft Bases H2O OH- F- R2S RSH RS-

CH3COO- PO43- SO4

2- I- SCN- S2O32-

CO32- ClO4

- NO3- CN- RNC CO R3P

ROH RO- R2O C2H4 C6H6 NH3 RNH2 N2H4 H- R-

Borderline C6H5NH2 C5H5N N3

- Cl- Br- NO2- SO3

2- N2 SCN- • Empirical classification of acids:

Class A: Hard Acids Class B: Soft Acids H+, Li+, Na+, K+ Cu+, Ag+, Au+, Tl+, Hg2+

Be2+, Mg2+, Ca2+, Sr2+, Mn2+ Pd2+, Cd2+, Pt2+, Hg2+, CH3Hg+, Co(CN)52-,

Pt4+, Te4+ Al3+, Sc3+, Ga3+, In3+, La3+, Tl3+, Tl(CH3)3, BH3, Ga(CH3)3

N3+, Cl3+, Gd3+, Lu3+ GaCl3, GaI3, InCl3 Cr3+, Co3+, Fe3+, As3+, CH3Sn+ RS+, RSe+, RTe+

Si4+, Ti4+, Zr4+, Th4+, U4+ Pu4+, Ce3+, Hf4+, WO4+, Sn4+

I+, Br+, HO+, RO+

UO22+, (CH3)2Sn2+, VO2+, MoO3+ I2, Br2, ICN, etc. Be(CH3)2, BF3, B(OR)3 trinitrobenzene etc. Al(CH3)3, AlCl3, AlH3 quinones

RPO2+, ROPO2+ TCNE

RSO2+, ROSO2

+, SO3 O, Cl, Br, I, N, RO•, RO2• I7+, I5+, Cl7+, Cr6+ M0 (metal atoms) RCO+, CO2, NC+ bulk metals

HX (hydrogen bonding molecules) CH2, carbenes Borderline

Fe2+, Co2+, Ni2+, Cu2+, Zn2+, Pb2+, Sn2+, Sb3+, Bi3+, Rh3+, Ir3+, B(CH3)3, SO2, NO+, Ru2+, Os2+, R3C+, C6H5

+, GaH3 • All important rule:

Hard acids prefer to bind to hard bases

Soft acids prefer to bind to soft bases

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E.g., Class A (hard) acids (such as Group 1 & 2 cations) bind to donor atoms of a base in the following order:

N >> P > As > Sb O >> S > Se ~ Te

F > Cl > Br > I

This order is essentially reversed if binding to a soft acid. • Consider the exchange reaction where B is some base:

The position of the equilibrium is a function of the hard-/softness of the base B:

Source: R.G. Pearson, J. Chem. Ed., 45, p. 581. N, O, F prefer to bind to the hard proton P, S, I, C prefer to bind to the soft CH3Hg+ RECALL: Notes on absolute hardness!

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Limitations of HSAB: • Consider the generic HOMO/LUMO diagram for hard/hard vs. soft/soft acid/base

interactions:

Source: Tarr & Miessler, Inorganic Chemistry, Prentice-Hall, 1998. The above MO suggests much higher stabilization for soft/soft! BUT: The MO diagram only considers covalent binding! • Important other contribution: Coulomb electrostatic forces

→ Relative adduct stabilities are difficult to predict based on HSAB alone.

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Other factors are: • Steric repulsions between acid/base

• Rearrangements of acid/base may be necessary for adduct formation, which will influence

the energetics (cf. BF3•NH3 experiment).

• HSAB is gas-phase only - what about the influence of solvents?

Quantitative prediction of adduct stabilities: The Drago-Wayland Equation → All limitations of HSAB are rooted in microscopic phenomena - difficult to measure. → Need a macroscopically measurable quantity to determine the stability of the adduct. Logical choice: Reaction enthalpy (measurable by calorimetry). Question: Can we find a single parameter (e.g. like pKas or redox potentials) that describes adduct strength as a f(ΔH)? A look at some experimental data:

ΔH [kcal/mol] Me3N Et2O Me2S Me3Ga 21 9.5 8

I2 12 4 8 → Me3Ga is a stronger acid than I2 in all cases BUT: Et2O is a stronger base than Me2S against Me3Ga and a weaker base than Me2S against I2 → order of basicity reversed as f(acid) ! Consequences: 1. There can be no reference point.

2. A single parameter approach will fail to explain the relative stabilities of Lewis acid/base

adducts. (… the problem involves more than one dimension …)

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Solution to this dilemma: Drago-Wayland Equation

−ΔH = EaEb + CaCb (ΔH is enthalpy of gasphase rxn: A + B → AB)

Lit.: R.S. Drago, J. Chem. Ed., 51, (1974), p. 300- 307.

• Two parameters for each acid and base: electrostatic (E) and covalent (C)

• ΔH is valid for gas-phase OR inert solvent.

• Experimental data for ΔH can be fit to this double parameterization. Drago arbitrarily set (in kcal/mol): I2: E = C = 1 DMA: E = 1.32 Et2S: C = 7.40 and measured 280 ΔH' s involving these species resulting in 144 variables. → 280 equations with 144 variables → solved for all E & C by least squares analysis over the entire data set (computer). Result: (Tables from Purcell + Kotz):

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Meaning and physical significance of E, C, and C/E values: • E, C, and C/E for any given single molecule are meaningless! • Can only compare Ea, Eb; Ca, Cb; Ca/Ea, Cb/Eb values of any two molecules.

• E and C can be interpreted as a measure of the ionic and covalent bond contributions,

respectively.

• E and C are NOT correlated with hardness and softness in any transparent manner. Applications: • With these numbers (tables) one can fairly accurately calculate and predict the ΔH's of

formation of over 1000 adducts.

• In total E and C values to calculate the ΔH's of about 37,000 adducts are known.

• Can calculate ΔH of adducts without doing the experiment.

• New donors/acceptors can be added to the list by measuring their ΔH with at least two other acids/bases (mathematical requirement: 2 equations - 2 variables). In reality more than two are usually done, in order to get better consistency.

• Can identify steric effects : e.g.: Me3N + BMe3 → Me3N-BMe3 Using E/C parameters: ΔH = 24.5 kcal/mol (derived using other adducts) Experimental data: ΔH = 17.6 kcal/mol (calorimetric) ΔΔH = 7 kcal/mol … all other things being equal this must be the result of steric repulsion ! • The last point should serve as a warning: The Drago-Wayland equation is supposed to be

quantitative, but actually - for any given two acid/base adducts - gets the relative stabilities right about 50 % of the time - use your discretion!

HOMEWORK: Exercises 4.19 – 4.27, 11.1 – 11.3, 11.6 – 11.11