Chemistry, Canadian Edition Chapter 19: Student Study Guide
Chapter 19: The Main Group Elements
Learning Objectives
Upon completion of this chapter you should be able to
• explain the chemistry of formation of adducts
• apply the concepts of hardness and softness to reactions
• explain the production, reactions, and uses of the main group metals
• explain the production, reactions, and uses of the metalloids
• explain the production, reactions, and uses of phosphorus
• explain the production, reactions, and uses of the other non-metals
Practical Aspects
The theme of this chapter is structure determines function. Many everyday uses for chemicals are explained based on their structures. This chapter introduces only two new themes: the concept of Lewis acids and bases and the concept of hard and soft acids and bases. This chapter then uses concepts from all previous chapters that emphasized structure, to explain the chemistry of the main group elements. Some examples are: drawing structures of covalent molecules, VSEPR theory, 3-D lattice arrays, band theory of solids, polymers, and redox and acid/base reactions.
19.1 Lewis acids and bases
Skills to Master:
· Identifying the Lewis acid and base in a reaction.
· Drawing the Lewis structure of an adduct.
Key Terms:
· Lewis acid – an electron pair acceptor.
· Lewis base – an electron pair donor.
· Adduct – product of a Lewis acid/base reaction: a Lewis base donates an electron pair to a Lewis acid, resulting in a new covalent bond formed between the two substances. In general, the formation of a Lewis acid-base adduct can be shown like this:
A + :B ® A–B
Key Concepts:
· A Lewis base must contain a pair of non-bonding electrons that are available for bond formation.
· A Lewis acid must be able to accept an electron pair, which means it either has vacant valence orbitals or is able to accommodate greater than 8 valence electrons. Most Lewis acids can be placed in one of these categories:
· A molecule with vacant valence orbitals (For example, BH3);
· A molecule with delocalized p bonds involving oxygen (For example, SO3);
· A metal cation (For example, Fe3+).
· The atoms involved in forming an acid-base adduct may need to re-hybridize in order for the adduct to be able to form.
Exercise 1: Identify each substance as a Lewis acid, a Lewis base, or neither. a) Co2+; b) CO; c) CH4; d) ethylenediamine; e) BeCl2; f) SnCl4.
Strategy: If the substance has an available (i.e., donatable) electron pair, then it is a Lewis base. If it can accommodate another electron pair, then it is a Lewis acid.
Solution:
a) Co2+: This is a metal cation, so it has space to accept an extra electron pair. It is a Lewis acid.
b) CO: Carbon monoxide has a triple bond between the C and O and a lone pair on each atom. It can act as a Lewis base.
c) CH4: Carbon is an n=2 element, so it cannot accommodate more than the 8 valence electrons it already has in this structure. It cannot act as a Lewis acid. There is no available lone pair, so it cannot act as a Lewis base. CH4 is neither.
d) ethylenediamine: recall that the structure of ethylendiamine contains two carbons singly bonded to each other, with an amine group attached to each carbon. Each nitrogen contains a lone pair, so ethylenediamine can act as a Lewis base.
e) BeCl2: Be has 2 valence electrons, each of which has formed a single bond with a Cl. Be is an n=2 element; it has space to accommodate 4 more electrons. This is a Lewis acid.
f) SnCl4: Sn in SnCl4 has used up all of its valence s and p orbitals, but since it is an n=5 element, it has available d orbitals, which can accept an electron pair, so it is a Lewis acid.
Try It #1: Which lone pair on carbon monoxide is a better Lewis base?
Exercise 2: Identify the Lewis acid and base in each reaction, indicate the formation of the adduct with curved arrows, and draw the 3-D structures of the products.
a) 6 H2O (l) + Ni2+ (aq) ® [Ni(H2O)6]2+ (aq); b) BF3 + F- ® BF4-;
c) SO2 (g) + CaO (s) ® CaSO3 (s)
Strategy and Solution: The Lewis acid will accept the lone pair that the Lewis base donates.
a) Ni2+ is the Lewis acid; H2O is the Lewis base. /b) BF3 is the Lewis acid; F- is the Lewis base. /
c) SO2 is the Lewis acid; O2- in CaO is the Lewis base. /
19.2 hard and soft lewis acids and bases
Skill to Master:
· Writing metathesis reactions.
Key Terms:
· Polarizability – the ease with which an atom’s electron cloud can be distorted by an electrical field (introduced in Chapter 8).
· Hard – term used to describe a substance whose electron cloud is not very polarizable.
· Soft – term used to describe a substance with a highly polarizable electron cloud.
· Hard-Soft Acid-Base (HSAB) Principle – hard Lewis acids tend to combine with hard Lewis bases; and soft Lewis acids tend to combine with soft Lewis bases.
· Metathesis reaction – reaction in which bonding partners are exchanged.
Key Concepts:
· Polarizability depends upon how strongly the valence electrons are attracted to the nucleus. Recall that an electron’s attraction to the nucleus depends upon distance (n value) and Zeff (amount of screening).
· The terms “hard” and “soft” are relative to each other. There is no absolute scale for assessing hard and soft characteristics.
Exercise 3: Rank each group from hardest to softest. a) Pd2+, Pt2+, Ni2+; b) Rb, In, Te; c) Co, Co2+, Co3+; d) Xe within: elemental Xe, XeF2, or XeF4.
Strategy: The more polarizable the electron cloud, the softer the substance.
Solution:
a) Pd2+, Pt2+, Ni2+: All of these elements are in the same column of the periodic table. All ions have the same charge, so we just need to compare the size of the electron clouds. Ni2+, with the smallest electron cloud, will be the hardest. The softest will be Pt2+, with the largest electron cloud.
b) Rb, In, Te: all of these elements have a valence shell of n=5, but in going from left to right across the periodic table, the Zeff increases. Te feels the pull of 15 more protons than Rb does, so Te will be the least polarizable, and the hardest. In will be the next hardest. Rb, with the lowest Zeff of the group, will have the most polarizable electron cloud and will be the softest.
c) Co, Co2+, Co3+: When an atom loses electrons, the remaining electrons can feel the attraction to the nucleus a little more strongly. Co3+, with the smallest radius, will be the hardest. Co2+, with the intermediate radius will be the next hardest. Co, with the largest radius, will be the softest.
d) Xe within: elemental Xe, XeF2, or XeF4? F is more electronegative than Xe, so it will pull electron density from Xe. The hardest will therefore be XeF4, then XeF2, and the softest will be Xe.
Try It #2: Which is harder, a) Sn2+ or Sn4+?
Exercise 4: Predict whether or not a metathesis reaction will occur. If a reaction does occur, predict the products.
a) TiI4 + 2 HgCl2 ® ? b) CaS + H2O ® ? c) CsF + LiI ® ?
Strategy: Hard acids prefer hard bases and soft acids prefer soft bases.
Solution:
a) Ti4+ is harder than Hg2+. Cl- is harder than I-. A metathesis reaction will therefore occur, to recombine the ions. The products will be: TiCl4 + 2 HgI2.
b) H+ is harder than Ca2+. O2- is harder than S2-. A reaction will not occur because the hard acid is already combined with the hard base, and the soft acid is combined with the soft base.
c) Li+ is harder than Cs+. F- is harder than I-. A reaction will occur to form: CsI + LiF.
Try It #3: Will BH3 prefer to react with NH3 or PH3?
19.3 The main group metals
Fast Facts – Aluminum:
· Main natural source: bauxite, Al(O)OH, contaminated with SiO2, Fe2O3, clay and other hydroxides.
· Isolation process: Convert bauxite to Al2O3, then use the Hall-Heroult process (an electrolysis reaction).
· Properties: lightweight, strong, and forms a single-layer oxide coating for protection.
· Uses: alloys are used for everything from aircraft bodies to beverage cans, AlCl3 is an important industrial catalyst, Al2(SO4)3 is used in water purification.
· Other:
· Third most abundant element in earth’s crust, and the most abundant metal.
· Stable as Al3+, so it’s difficult to form Al (s).
· Hard Lewis acid.
Fast Facts – Lead:
· Main natural source: galena, PbS
· Isolation process: heat with oxygen to form PbO, then reduce with charcoal (C).
· Uses: lead storage batteries (for cars); alloys such as pewter.
· Other:
· Stable in +2 and +4 oxidation states.
· Highly toxic, so its present-day use is limited. For many years, it was used in paints, solder, and as an anti-knock additive in cars.
Fast Facts – Tin:
· Main natural source: cassiterite, SnO2
· Isolation process: add charcoal (C, a good reducing agent) at high temperature.
· Properties: relatively low melting point, resists corrosion.
· Uses: alloys such as pewter, bronze, and solder; coating for aluminum cans.
· Other:
· Stable in +2 and +4 oxidation states.
Exercise 5: Explain why, in nature, Sn and Al exist as oxides while Pb exists as a sulphide.
Strategy: Compare cations and anions to see if there is a trend.
Solution: Oxygen and sulphur are both in the same group; both oxide and sulphide have the same charge (-2). The difference between the two is that S is larger than O, making S2- a softer base than O2-. Similarly, Pb2+ is larger and softer than Sn4+ or Al3+. Cations with high charges like +3 or +4 are typically considered hard. It seems reasonable, then, that Pb2+ will be more attracted to S2- , while Sn4+ and Al3+ will be more attracted to O2-.
19.4 the metalloids
Skills to Master:
· Explaining the bonding in boron compounds.
Fast Facts – Silicon:
· Main natural source: silicon dioxide, SiO2, and related silicate anions.
· Isolation process: to obtain up to 98% purity, heat with coke (charcoal) at extremely high temperature. A multi-step process is required for high-purity silicon for semiconductors.
· Properties: semiconductor.
· Uses: computer chips, glass, and silicone polymers.
· Other:
· Second-most abundant element in earth’s crust.
Fast Facts – Other Metalloids:
· B – unique properties due to its small size: high ionization energy, has no available d orbitals.
· Sb – alloyed with lead in a lead storage battery to minimize chance of water undergoing electrolysis during battery recharging.
· Ge – semiconductor.
· As – pesticide.
· Binary compounds of some metalloids, such as GaAs or InSb, have the same number of valence electrons as Si or Ge, and exhibit semiconductor properties.
Exercise 6: Compare the reduction potentials of Sb and water to determine how the presence of Sb in a lead storage battery can minimize the chance of water electrolysis during battery recharging. Given:
Sb + 3 H+ (aq) + 3 e- « SbH3, E° = -0.150 V.
Strategy: Look up the standard reduction potential for the reduction of water in Appendix F of the text and compare its value to that of the antimony half-reaction’s reduction potential. The substance with the greatest reduction potential will get reduced.
Solution:
Half-Reaction
/Standard Reduction Potential
E° (V)Sb + 3 H+ (aq) + 3 e- « SbH3 / -0.510
2 H2O (l) + 2 e- « H2 (g) + OH-(aq) / -0.828
The reduction of Sb (in an acidic environment) and the reduction of water are both non-spontaneous processes. When the battery is recharged, electrical potential is put into the battery to drive the non-spontaneous processes in the forward direction. The Sb half-reaction has a higher reduction potential than the water half-reaction, so Sb will preferentially undergo reduction.
19.5 phosphorus
Skills to Master:
· Describing the bonding in red and white phosphorus.
· Explaining the processes for making phosphoric acid and phosphorus fertilizers.
· Explaining phosphate condensation reactions.
Fast Facts – Phosphorus:
· Main natural source: Apatite, Ca5(PO4)3X, where X = F, OH, or Cl.
· Isolation process: heated with silica and coke (C, a good reducing agent) at high temperature, in the absence of oxygen.
· Properties: Phosphorus exists in three forms:
· White phosphorus (P4) – individual small molecules; highly reactive and highly toxic.
· Red phosphorus – polymer made from heating white phosphorus. Less reactive and less toxic than white phosphorus.
· Black phosphorus – contains long, crosslinked chains of P atoms. More stable than red phosphorus.
· Uses: phosphoric acid (one of the top 10 chemicals produced in the U.S.), phosphate fertilizers, polyphosphate detergents, herbicides, insecticides, and biological molecules like ATP.
· Other:
· Almost 90% of phosphoric acid produced goes towards making phosphate fertilizers.
· The most important phosphate fertilizer is (NH4)2HPO4.
· Elevated concentrations of phosphates in the environment (primarily from detergents), cause excess algae growth in lakes, which deplete the lakes of oxygen, killing off aquatic life.