CHEM1910 Inorganic Lecture 1 Introduction PDF

CHEM1910 Block 1: Chemistry of the Main Group Elements January 23rd 2024 Dr Nickeisha Stephenson  CHEM1910 – Introductory Chemistry III This course is divided into three blocks: I. Main Group Chemistry II. Transition Metal Chemistry III. Molecular Energetics Bui...

CHEM1910 Block 1: Chemistry of the Main Group Elements January 23rd 2024 Dr Nickeisha Stephenson  CHEM1910 – Introductory Chemistry III This course is divided into three blocks: I. Main Group Chemistry II. Transition Metal Chemistry III. Molecular Energetics Builds on the topics covered in CHEM1810 and CHEM1820 The course is assessed by your performance on three graded tutorials (online), two 1- hour course tests, face-to-face), and a two-hour exam (face-to-face). I will teach the first two blocks and Dr Mark Lawrence will teach the final block.  CHEM1910 - LECTURE & TUTORIAL SCHEDULE 2023/2024 Academic Week 1 2 3 4 5 6 7 8 9 10 11 12 13 Beginning Monday 22-Jan 29-Jan 5-Feb 12-Feb 19-Feb 26-Feb 4-Mar 11-Mar 18-Mar 25-Mar 1-Apr 8-Apr 15-Apr Tue 8 am/ Tue 5 pm MG MG MG MG TM TM TM TM MSE MSE MSE MSE No Class Thu 12 pm / Fri 5 pm MG MG MG MG TM TM TM MSE MSE MSE MSE Test Day Course Test 1 – Week 5 Thursday, February 22nd at 5 PM MG = Main Group Chemistry TM = Transition Metal Chemistry Course Test 2 – Week 10 Thursday, March 28th at 5 PM MSE = Molecular Spectroscopy and Energetics Required Course Textbook: Chapters 25 – 28 Lecture Topic 1 Course overview, Periodic trends review 2 Hydrogen and its Compounds (Chapter 25) 3 S-Block Chemistry – Group 1 (Chapter 26) 4 S-Block Chemistry - Group 2 (Chapter 26) 5 P-block Chemistry – Groups 13, 14, and 15 (Chapter 27) 6 P-Block Chemistry – Groups 16 (Chapter 27) 7 P-block Chemistry –Groups 17 and 18 (Chapter 27) d-block chemistry, Coordination chemistry, geometry, and 8 and 9 calculating oxidation state (Chapter 28) 10 Transition metal complexes – Isomers (Chapter 28) 11 and 12 Introduction to Crystal Field Theory (Chapter 28) Colour and Magnetic Properties of transition metal 13 complexes (Chapter 28) 14 Stability of transition metal complexes (Chapter 28)  Engage with your textbook, class material, tutorials etc. Ask questions – If you have a question chances General Rules are others around you do too. Speak up. This a judgment free zone. and Expectations No pictures or video in class without my permission. If I catch you, I will ask you to leave. Keep phones on silent.  Lecture slides will be made available on OurVLE on Tuesday and Thursday by the end of the day. General Rules Go to tutorials – Tutorials will feature the questions I will likely ask you on an exam. and Expectations If you have difficulties grasping the material, seek help – ASU, Lecturer, demonstrator, etc. Cont’d Avoid marathons!!! Deliberate practice. Questions Before We Get Started? Chemistry of the Main Group Elements  The Four Major Groups of the Periodic Table 1. 1. 2. 3. 4.  The Importance of Main Group Elements Group 1 – alkali Metals Group 2 – alkaline earth metals Group 13 – no trivial names Group 14 - no trivial names Group15– pnictogens Group 16 – chalcogens Group 17 – halogens Group 18 – noble gases The 10 most abundant elements by mass (a) in the earth's crust and (b) in the human body. All are main-group elements except iron and titanium.  Main Group Chemicals in Industry – Top 10 Sulfuric acid (H2SO4) 47.7mil tons - Fertilizers, chemicals, oil refining Nitrogen (N2) 40.2 mil tons - Inert atmospheres, low temperatures, ammonia production Oxygen (O2) 29.8 mil tons - Steelmaking, welding, medical uses Ethylene (CH2=CH2) 24.5 mil tons - Plastics, antifreeze Lime (CaO) 20.9 mil tons - Steelmaking, chemicals, water treatment Ammonia (NH3) 19.5 mil tons - Fertilizers, nitric acid Phosphoric acid (H3PO4) 13.1 mil tons – Fertilizers, detergents Chlorine (Cl2) 12.6 mil tons – Chemicals, plastics, water treatment Propylene (CH3CH=CH2) 12.6 mil tons – Plastics, fibres, solvents Sodium hydroxide (NaOH) 11.8 mil tons - Chemicals, textiles, soaps The Importance and Utility of the Periodic Table in our study of the elements Ø Physical and Chemical properties of elements and their compounds follow largely predictable patterns down a group in the periodic table Ø Physical and Chemical properties of elements and their compounds follow largely predictable patterns across periods in the periodic in the periodic table Ø This is called periodicity, and it can be summarized in the periodic law Ø Periodic Law : The properties of an element are periodic functions of their atomic numbers The Big Idea: The structures of atoms determine their properties, and hence the behaviour of elements is related to where they are located on the periodic table. Ø The chemical reactivity of the elements is largely determined by their outermost or valence electrons – Recall Zeff Ø The similarity of the valence electron configuration is what makes elements in the same group have similar chemical properties. Ø Electron configuration controls the number of bonds that an element can form and will affect its chemical and physical properties. Ø As we saw in CHEM1810, five atomic properties are largely responsible for the characteristic properties of each element Ionisation energy Electron affinity Atomic radii Effective Nuclear Charge, Zeff Electronegativity Polarizability Periodic Trends  Effective Nuclear Charge – Zeff The nuclear charge felt by an electron when both the actual nuclear charge (Z) and the repulsive effects (shielding) of the other electrons are considered. Ø Increases across a period. Nuclear charge increases with no significant increase in shielding. n is constant Ø Decreases down a group. The nuclear charge increases down a group, However, the successive addition of core electrons increases shielding effect and counters this effect. n increases Ø Zeff is behind all other periodic trends Recall Slater’s rule and the calculation of Zeff Atomic Radii  Ionization Energy The first ionisation energy (IE1) is the energy required to remove the first valence electron from a mole of gaseous atoms  Ionization Energy The first ionisation energy (IE1) is the energy required to remove the first valence electron from a mole of gaseous atoms e Decreased number of electrons + increased Zeff = smaller radius 21 Electron Affinity Vs. Electron gain energy e Ø The energy released when a neutral atom gains an electron to form an anion. - IUPAC definition of electron affinity Ø The change in energy associated with the addition of an electron to an atom in the gaseous state. - Electron gain energy Negative Electron Gain Energy = Favourable Reaction – Exothermic In this example, addition of an electron to fluorine atoms, results in a product, fluoride anion that is lower in energy F (g) + e- than fluorine atom. This reaction is exothermic, energy is released and results in a F 1-(g) product more stable than the starting the starting material. This reaction is favorable and will want to proceed. Negative change in energy. Positive Electron Gain Energy = Unfavourable Reaction – Endothermic In this example, addition of an electron to magnesium atoms, results in a product, Mg anion that is higher in energy than Mg 1-(g) magnesium atom. This reaction is endothermic, requires energy and results in a product less stable than the starting the starting material. This reaction is unfavorable Mg (g) + e- and will not want to happen. Positive change in energy  Recall exceptions to the trend - Groups 2, 15, and 18 Trend in Electron Gain Energy Electronegativity Electronegativity No difference Moderate in E/N difference in E/N 27 Electrons Electrons shared shared equally unequally Electronegativity (EN) The measured tendency of an atom to attract a bonding pair of electrons. Electronegativity and Chemical Bonds The difference in EN of the atoms involved determines the nature of the chemical bond formed The relationship between the ionic character of a covalent bond and the electronegativity difference of the bonded atom.  Polarizability Ø A measure of the ease with which an electron cloud can be distorted. It is greatest for electron-rich heavier atoms of a group and negatively charged ions. Ø Polarizing power is the ability to distort the electron cloud of a neighbouring atom or ion. Atoms or ions with small size and high positive charge are described as having high polarizing power. Ø Polarizability and polarizing power help to explain why aluminium chloride exists as a covalent molecule rather than an ionic solid Review 1. How would you expect the atomic radius to change in the series B, C and N? A. B>C>N B. N>C>B C. C>B>N D. NB 2. Put the following elements in decreasing order of first ionisation energy A. C, N, O B. N, O, C C. C, O, N D. N, C, O 3. Which one of the following statements is not true about H2? A. The H-H bond is covalent B. The H-H bond is polar C. The molecule can be described as a homonuclear diatomic D. The electron pair is shared equally between atoms Questions taken from Chemistry3 student resources website 4. Which of the following elements has atoms with the greatest polarizability? A. Carbon B. Aluminium C. Fluorine D. Germanium 5. An element “E” in period four forms a molecular hydride with the formula HE. What is the identity of E? 6. Which of the following elements has the greatest electron affinity? A. Nitrogen B. Oxygen C. Gallium D. Tellurium  7. For which of the following molecules does octet theory work? A. PF5 B. NO C. [PtCl2(NH3)2] D. N2 8. According to Pauling electronegativity theory, which one of the following species would be the most polar? A. HCl B. HI C. HF D. HBr 9. When two s atomic orbitals combine, which of the following bond orders is not possible from this combination whatever the filling of the corresponding orbitals. A. 1 B. 0.5 C. 0 D. 2 Questions taken from Chemistry3 student resources website  10. How would you expect the bond strength to change in the series C2 to C2- to C22- A. Get stronger then weaker B. Get weaker C. Get stronger D. Get weaker then stronger 11. Taking z as the internuclear axis and assuming that the energies of the orbitals are similar, which of the following pairs of orbitals would result in a non-bonding combination? A. 2pz and 2s B. 2py and 2s C. 2s and 2s D. 2pz and 2pz For answers and additional questions see quiz on OurVLE Questions taken from Chemistry3 student resources website  Bond order of 0 indicates that Bond order of 1 corresponds to He2 is not stable at normal H2 having a single bond conditions.  s*2pz s2pz p*2px , p*2py p2px , p2py s2pz s*2ps s2s In general the energies of the orbitals decreases across the period as the effective nuclear charge increases and atomic radius decreases. Between N2 and O2 the order of the orbitals changes. n=2 level where p orbitals exist then you must show the MOs formed from them even if orbitals not filled.  CHEM1910 Block 1: Chemistry of the Main Group Elements January 25th 2024 Dr Nickeisha Stephenson  Hydrogen A Unique Element* *So much so that your textbook has devoted an entire chapter to it –Chapter 25! – Read and discover Where we are going this Lecture…. o We will discuss the element hydrogen and its placement on the periodic table. o We will learn about some of the physical properties of hydrogen and its isotopes. o List some of its uses. o Discover how hydrogen is made on an industrial scale, using electrolysis and from the reactions of dilute acids with metals. o Discuss the hydrogen economy o Learn about hydrogen and its compounds by investigating the various compounds hydrogen forms with other elements in periods 2 and 3. o Revisit hydrogen bonding. Resources Chapter 25 of Chemsitry3 (posted on OurVLE) Notes from Prof Lancashire - http://wwwchem.uwimona.edu.jm/courses/CHEM19 02/IC10K_MG_hydrogen.html Where to place Hydrogen on the Periodic Table?  Hydrogen – 1s1 Where does hydrogen belong on the periodic table ? With the halogens or with alkali metals? Hydrogen does not belong in group 17  Hydrogen: ü Gains an electron to form a monoanion (H-) ü Monoanion has a noble gas configuration (1s2) ü A gas at room temperature ü Forms homonuclear diatomic molecules (H2) ✗ Forms a cation ✗ Reactivity is very different from the halides Where does hydrogen belong on the periodic table? Group 17? Placed in group 17 because the addition of one electron gives a noble gas configuration  Where does hydrogen belong on the periodic table? Group 1? Hydrogen: ü Electron configuration is 1s1 Placed in group 1 because of electronic configuration (1s1) ü Forms a monocation like group 1 elements (H+ ) ✗ Unlike group 1 elements, monoanion H- ions are common ✗ A gas at standard room temperature and pressure ✗ Forms homonuclear diatomic molecules (H2) ✗ Not as reactive as group 1 elements ✗ In general, its chemistry does not match up with the members of the group. Comparing hydrogen's atomic properties with those of Li and F greater demonstrates the issue. Hydrogen has properties that resemble both of these elements. For this and other reasons, hydrogen is a truly unique element. Hydrogen does not fit neatly in the Periodic Table. Often placed in the center of the periodic table. Where we are going this Lecture…. ü We will discuss the element hydrogen and its placement on the periodic table. Ø We will learn about some of the physical properties of hydrogen and its isotopes. o List some of its uses. o Discover how hydrogen is made on an industrial scale, using electrolysis and from the reactions of dilute acids with metals. o Discuss the hydrogen economy o Learn about hydrogen and its compounds by investigating the various compounds hydrogen forms with other elements in periods 2 and 3. o Revisit hydrogen bonding.  Hydrogen is the most abundant element in the universe, accounting for about 75% of the universe's mass. On Earth, hydrogen is rarely found in its pure state and is often found combined with oxygen (water) or carbon (hydrocarbons). Dihydrogen is very light and escapes Earth’s gravity - 0.53 ppm H2 by volume in the atmosphere. Hydrogen forms compounds and molecules with many transition elements and with every main group element except the noble gases, indium, and thallium. Hydrogen & Dihydrogen The earth’s crust hydrogen is the 9th most abundant element on a mass basis (0.9%) and 3rd most abundant on an atom basis (15.4%) Industrially large amounts of hydrogen are produced for the synthesis of chemicals such as ammonia (NH3) and methanol (CH3OH). Dihydrogen is a colourless, odourless, and tasteless gas made up of diatomic H2 molecules. Dihydrogen is relatively unreactive at room temperature and at elevated temperatures as homolysis of the H–H requires large amounts of energy. H2(g) → 2H(g); ΔH = + 436 kJ/mol  Hydrogen Has Nuclear Spin Isomers In dihydrogen, the two electrons in the molecule will be spin paired, but there is no similar requirement for the two nuclei; they may be parallel or opposed. This results in two nuclear spin isomers which are called ortho and para-dihydrogen. Recall that electrons have a property called spin that describes their angular momentum. Nuclei of atoms also have angular momentum and, therefore, have nuclear spin. Parahydrogen has antiparallel nuclear spins. The proton of the hydrogen atom has a spin of +1/2 or -1/2. Orthohydrogen has parallel nuclear spins Not all nuclei have the same spin property. At standard temperature and pressure, hydrogen gas Nuclear spin varies by element. contains about 25% of the para form and 75% of the ortho form  Hydrogen Has Nuclear Spin Isomers In dihydrogen, the two electrons in the molecule will be spin paired, but there is no similar requirement for the two nuclei; they may be parallel or opposed. This results in two nuclear spin isomers which are called ortho and para-dihydrogen. Parahydrogen has antiparallel nuclear spins. Orthohydrogen has parallel nuclear spins At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the When cooled to 20K, orthohydrogen converts to almost ortho form pure parahydrogen.  Hydrogen has Three Isotopes Isotopes – Elements that have the same number of protons but differ by the number of neutrons. Symbol: H Symbol: D Symbol: T Standard Two times as Three times as hydrogen heavy as protium heavy as protium  Hydrogen Isotopes Abundance: 1 H (protium H, 99.985 %) 2H (deuterium D, 0.015 atom % ) 3 H (tritium T, ~ 10-16 atom % ; radioactive, t ½ = 12 years) Chemical Properties of the isotopes are similar as chemical behaviour is determined mainly by electronic structure – (1s1) OH The difference in mass between isotopes is usually small and does not lead N to significant differences in properties. Me N 3H Me 3C Deuterium and tritium are double and triple the mass of protium, and this leads to a quantitative difference in physical properties: Visualization of muscarinic - Melting point receptors in the brain using a - boiling points tritium containing molecule - The density - Rates of reactions (kinetic isotope effect) Physical Properties of Hydrogen Isotopes Some physical properties of the hydrogen isotopes. Interatomic isotope MP /K BP /K ΔHdiss /kJmol-1 Distance /pm H2 13.99 20.27 435.99 74.14 D2 18.73 23.67 443.4 74.14 T2 20.62 25.04 446.9 74.14  H2O (s) D2O (s) A glass of H2O (l) with ice and ‘heavy’ ice Comparison of H2O and D2O Rats can distinguish D2O from Property H2O D2O H2O; given the choice, they will Melting point /K 273.15 276.97 choose H2O. Boiling point /K 373.15 374.5 Replacing 50% of an organism's Temperature of maximum 277 284.2 density /K water with D2O can result in death Maximum density /g cm3 0.99995 1.1053 Due to the differences in boiling point, D2O can be separated from H2O by fractional distillation. Standing bodies of water with high rates of evaporation becomes enriched in D2O over time. Isotopic Differences https://www.youtube.com/watch?v=YYInVraBe7s  Where we are going this Lecture…. üWe will discuss the element hydrogen and its placement on the periodic table. ü We will learn about some of the physical properties of hydrogen and its isotopes. Ø List some of its uses. Ø Discover how hydrogen is made on an industrial scale, using electrolysis and from the reactions of dilute acids with metals. o Discuss the hydrogen economy o Learn about hydrogen and its compounds by investigating the various compounds hydrogen forms with other elements in periods 2 and 3. o Revisit hydrogen bonding. Preparation of H2 Preparation of Hydrogen – Global Production Ø 48 % of the global demand for H2 is generated from natural gas Ø 30% from oil/ naphtha Ø 18% coal Ø 3.9 % from electrolysis Hydrogen does not occur native on earth as H2 gas Preparation of Hydrogen Industrial Production Steam Hydrocarbon re-Forming Process Step 1 Step 2 Step 3 - Most utilised industrial process for producing H2 - Three-step process  Steam Hydrocarbon Reforming Process Step 1: Reformer: H2O (g) + CH4 (g) 1100 ºC CO (g) + 3 H2 (g) ΔH = + 206 kJ/ mol Ni cat. Synthesis “syn” gas The combination of steam and methane to make synthesis gas Step 2: Water-gas shift reaction: 400 ºC CO (g) + H2O CO2 (g) + H2 (g) ΔH = -41 kJ/ mol cat. Shifts the composition of synthesis gas by removing toxic CO and producing more H2  Steam Hydrocarbon Reforming Process (Cont’d) Step 3: Purification CO2 (g) + 2 OH- (aq) CO3 2- (aq) + H2O (l) H2 / CO2 mixture is passed through a basic solution that converts CO2 to carbonate o Purity of H2 à 97-98% o Method produces CO2 as a by-product o The largest consumer of hydrogen is the Haber process Preparation of Hydrogen – Electrolysis (small-scale production) Water oxidation half reaction (anode): 2H2O → O2 + 4H+ + 4e- Reduction half reaction (cathode): 4H+ + 4e- → 2H2 Overall water splitting reaction: 2H2O → O2+ 2H2 Ø H2 of greater than 99.95 % purity is made by the electrolysis of water Ø Involves the breaking of strong O-H bonds Ø Requires a large amount of energy, 286 kJ per mole of H2 produced. Not practical for large-scale production unless electricity is cheap. Ø Can be used to separate deuterium from protium electrolysis 2 H2O (l) 2 H2 (g) + O2 (g) Faster Kinetic Isotope effect electrolysis 2 D2O (l) 2 D2 (g) + O2 (g) Slower Preparation of Hydrogen – Action of dilute acid on active metals Ø Redox reaction between Zn metal and aqueous acid to produce H2(g) and Zn2+salt solution Ø laboratory-scale production Preparation of Hydrogen – Action of dilute acid on active metals Eooverall = Eoox + Eored Eored = 0.00 V Eoox = 0.76 V Zn + 2 H+ Zn2+ + H2 Eo = Eoox + Eored = 0.76 V Preparation of Hydrogen – Action of dilute acid on active metals In principle any metal with a negative standard reduction potential will react with an acid to generate H2 How hydrogen was discovered in 1671 and rediscovered in 1766  The Hydrogen Economy Ø On going research into making the electrolysis/ decomposition of water to produce H2 a viable industrial process; the vision of hydrogen economy Ø Heat liberated when H2 burns is more than twice that of gasoline, oil or natural gas on a mass basis Ø Electrolysis of water offers a way to produce H2 without producing CO2 Photoelectrochemical Cell Renewable, but intermittent energy source hn Anode Cathode Oxygen Evolution Hydrogen Evolution Catalyst Catalyst h+ e- 2H2 (Fuel) 2H2O semiconductor to absorb light Abundant feedstock (starting material) O2 4H+ 2001 rate of energy consumption of 13.5 TW (86% from fossil fuels) 2050 consumption rate is projected to be 27 TW Solar energy is abundant and can yield carbon-neutral fuels Lewis, N.S.; Nocera, D.G. Proc. Natl. Acad. Sci. 2006, 103, 15729. Interested to learn more about the downsides of the hydrogen economy? https://www.youtube.com/watch?v=Zklo4Z1SqkE  Reactivity of Hydrogen 1. Loss of electron to form H+ (similar to alkali metals) H(g) → H+(g) + e- Ei = +1312 kJ/mol 2. Gain an electron to form H- (similar to halogens) H(g) + e- →H-(g) Eea = -73 kJ/mol 3. Formation of an electron-pair bond in covalent molecules (H·, hydrogen atom ) CH4, NH3, H2O, HF (neutral H compounds)  Reactivity of Hydrogen Ø Large ionization energy, hydrogen doesn’t completely transfer its electron in chemical reactions. Shares electrons with a nonmetallic element to form covalent compounds ØIn liquids and solids a bare proton is too reactive to exist by itself, instead it bonds to a molecule that has a lone pair of electrons e.g. in H2O à H3O+ ØAdding an electron to H releases energy (less than the halogens). Hydrogen will accept an electron from an active metal to give an ionic hydride e.g. NaH or CaH2  Compounds of Hydrogen- Binary Hydrides Hydrogen forms binary compounds with most elements of the Periodic Table. Hydrides have a wide variety of properties and structures. These compounds are called hydrides. Any compound that consists of hydrogen and one other element is known as a binary hydride. The key to understanding the differences in the chemistry of these hydrides lies in the electronegativity of hydrogen. For Hydrogen, cP = 2.20  Binary Hydrides of Period Two BeH2 LiH B2H6 CH4 NH3 H2O HF ,..... Most Common hydrides of the s- and p- block 25.2 COMPOUNDS OF HYDROGEN Electron-deficient hydride: not possible to describe the bonding of the structure in terms Group 1 2 3-12 13 14 15 16 17 of 2-centre 2-electon bonds (e.g. BeH2 and BeH2 B2H6. '''?~ ' CH 4 I ··. ,_,. N'::1 3 ~ 20 I ' HF group 13 hydrides. beryllium diborane methane ammonia~ water hydride. · ~,ii,. , ,·~, q ·""- \i ~~JJ.~tK~ J , ,,,,::, ,;: Electron –precise: all the valence electrons of AIH 3 aluminium the central atom are involved in forming bonds hydride (grp 14) Sc-Zn Electron-rich: Not all the electrons on the central atom participates in bonding. Grps 15- Y-Cd 17 Ionic hydrides Electron-deficient covalent hydrides La-Hg Electron-precise covalent hydrides """' Electron-rich covalent hydrides Hydride is not we\\ characterized Figure 25.1 The most common hydrides of the s- and p-block elements. Where are the electrons? c and bonding in binary hydrides  CHEM1910 Block 1: Chemistry of the Main Group Elements Lecture 3 January 30th 2024 Dr Nickeisha Stephenson Compounds of Hydrogen – Binary Hydrides Hydrogen forms binary compounds with most elements of the Periodic Table. Hydrides have a wide variety of properties and structures. These compounds are called hydrides. Any compound that consists of hydrogen and one other element is known as a binary hydride. The key to understanding the differences in the chemistry of these hydrides lies in the electronegativity of hydrogen. For Hydrogen, cP = 2.20 Binary Hydrides of Period Two BeH2 LiH B2H6 CH4 NH3 H2O HF Most Common hydrides of the s- and p- block The Standard Gibbs Energy of Formation (∆fG°) is a thermodynamic quantity that represents the change in Gibbs free energy when one mole of a compound is formed from its elements in their standard states at a specified temperature (usually 25°C or 298 K). It is expressed in units of kilojoules per mole (kJ/mol). When examining s- and p-block hydrides, the ∆fG° values provide insights into the thermodynamic favorability of their formation. A negative ∆fG° value indicates that the formation of the hydride, from hydrogen and its constituent elements is exergonic, meaning it releases energy and is thermodynamically favoured under standard conditions. This suggests that the hydride is stable relative to its constituent elements at room temperature. Shriver and Atkins 5th edition The s-block elements form thermodynamically stable hydrides (e.g., NaH, CaH2), with negative Gibbs free energies of formation (∆fG° < 0), indicating their stability relative to their constituent elements at room temperature. The formation of an ionic crystal lattice compensates for the energy required to cleave the H2 bond. For the p-block elements, there is a variable trend in the stability of hydrides. Notably, the first members of each group (CH4, NH3, H2O, HF) form stable hydrides, but this stability decreases for heavier congeners in each group. This trend is illustrated by the decreasing E–H bond energies as we go down the group. Influence of Atomic Properties on Bond Strength: The strength and stability of hydrides are influenced by atomic properties. Strong E-H bonds are required for a compound to be thermodynamically stable. In p-block hydrides, bond strength is highest with Period 2 elements and decreases down the group, attributed to poor orbital overlap between the compact 1s orbital of hydrogen and the more diffuse s and p orbitals of heavier p-block elements. For a binary hydride to be exergonic and stable with respect to its elements it need to have E–H bonds that are stronger than H–H bond Where are the electrons? c and bonding in binary hydrides  Where are the Electrons? A Thought Experiment A B Ø If the bonding pair of electrons is shared equally between A and B, what properties besides electronegativity could we assign to these elements? Where are the Electrons? A Thought Experiment A B Ø If the bonding pair of electrons is shared equally between A and B, what properties besides electronegativity could we assign to these elements? Element A is equal to Element B Element A and Element B have similar Zeff Where are the Electrons? A Thought Experiment A B Ø If the bonding pair of electrons is shared unequally between A and B, with B “hogging” the electron pair, what properties besides electronegativity could we assign to these elements? Where are the Electrons? A Thought Experiment A B Ø If the bonding pair of electrons is shared unequally between A and B, with B “hogging” the electron pair, what properties besides electronegativity could we assign to these elements? Element A Element B Lower Zeff than B Higher Zeff than A Lower ionization energy than B Higher ionization energy than A Larger atomic radii than B Smaller atomic radii than A More positive electron affinity*** More negative electron affinity*** Maybe a metal Maybe a nonmetal  Where are the Electrons – Electronegativity Electronegativity (c) is a measure of the tendency of an atom to attract electrons (electron density) towards itself. The more strongly an atom attracts bonding electrons towards itself, the larger its electronegativity. Electronegativity is a concept, not a physical property and can not be measured directly. Rather it is calculated using various pieces of experimental data. Ø Ionization energy and electron affinity, Mulliken Scale, cM (1934) Ø Effective nuclear charge and covalent radii, Allred-Rochow scale, cAR (1958) Ø Bond enthalpy, Pauling Scale, cP (1932)  Mulliken’s Electronegativity Scale, cM Robert S. Mulliken observed that elements with high ionization energy (the electrons are not readily released) and a high electron affinity (it is energetically favourable to acquire electrons) lead to a high electronegativity (c). Mulliken's scale relies on ionization energy and electron affinity measured in electron volt (eV): An element with small ionization energy and a small | E + Eea | 𝜒! = I1 positive value for electron affinity will have a lower 2 value in the numerator and hence, a low electronegativity. Using Mulliken’s method, the electronegativity of fluorine is 1004.6 kJmol-1. Mulliken’s electronegativity values can be converted to Pauling’s by dividing by 252.4 kJmol-1 this gives a value of 3.98 for fluorine; highlighting that even though both methods rely on different measurements they seem to be measuring the same phenomenon. Method is limited as electron affinity data is not readily available for many elements. Allred-Rochow Electronegativity Scale, 𝜒 !" Given that elements with high electronegativity values have small atomic radii and large effective nuclear charge, scientists Allred and Rochow used these values to produce their electronegativity scale. 𝜒 "# is based on the coulombic force of attraction of the atom for electron density in the bond, Zeff values based on Slater’s rule as, such the 𝜒 "# scale is not very rigid. Factors were added to make the scale more similar to Pauling’s. Can be used to measure the electronegativity of the noble gases. Method is limited by the availability of covalent radius data. Electronegativity: The Pauling's Scale Ø Pauling introduced the concept of electronegativity, and the Pauling scale is the most widely accepted scale of electronegativity (1932). Linus Pauling (1901-1994) Ø Scale relies on bond enthalpies (D in eV) Winner of the Nobel prize Ø Pauling’s model considered a hypothetical molecule AB and In chemistry and Peace postulated that the electronegativity of atoms A and B could be determined by comparing the measured bond energy of the A–B molecule with the expected (calculated) bond energy of the A–B molecule. Ø The difference (DD) between the actual (measured) and the expected (calculated) bond enthalpy can be used to determine cP of atoms A and B.  Electronegativity: The Pauling's Scale Ø The difference in electronegativities between A and B: #" 𝜒! 𝐴 − 𝜒! 𝐵 = (∆𝐷) $ When using eV 𝜒! 𝐴 − 𝜒! 𝐵 = 0.102(∆𝐷) ⁄" ! When using kJmol-1 Ø DD is the difference between the measured value of bond enthalpies for molecule AB and the mean bond enthalpy for A2 and B2 (expected bond enthalpy) 𝐷 &–& + 𝐷 (–( ∆𝐷 = 𝐷 &–( − 2 Ø Pauling assigned the electronegativity of fluorine as 4.00 and this allowed the electronegativity of all other elements to be determined. See Chemistry3 Chapter 4, page 177. Using Pauling’s Scale to calculate cp(H) H2 (g) 2 H (g) , D = +435.8 kJmol-1 F2 (g) 2 F (g) , D = +158.7 kJmol-1 H—F H (g) + F (g), D = +569.7 kJmol-1 )( 𝐷 $–$ +𝐷 &–& 𝜒' 𝐹 − 𝜒' 𝐻 = 0.102(∆𝐷) * ∆𝐷 = 𝐷 $–& − 2 | 4.00 – 𝜒 ' 𝐻 | = 0.102(272.4)1/2 435.8+158.7 = 569.7 − 2 𝜒 ' 𝐻 = 2.3 = 272.4  Pauling’s Scale: Bonding Between Unlike Atoms H2 (g) 2 H (g) , D = +436 kJmol-1 Average bond dissociation enthalpy H-I (expected) = I2 (g) 2 I (g) , D = +151 kJmol-1 (436 kJmol-1 + 151 kJmol-1 ) / 2 = 294 kJmol-1 H—I H (g) + I (g), D = +298 kJmol-1 D = (298 kJmol-1 – 294 kJmol-1 ) = 4 kJmol-1 H2 (g) 2 H (g) , D = +436 kJmol-1 Average bond dissociation enthalpy H-F (expected) = F2 (g) 2 F (g) , D = +159 kJmol-1 (436 kJmol-1 + 159 kJmol-1 ) / 2 = 297 kJmol-1 H—F H (g) + F (g), D = +570 kJmol-1 D = (570 kJmol-1 – 297 kJmol-1 ) = 273 kJmol-1 The mean calculated bond energy for H-I is similar to the experimentally determined, measured bond energy for the H-I molecule. This indicates that DEN between H and I is small. The mean calculated bond energy for H-F is lower than the measured bond energy. Pauling saw this as evidence that H and F were ”unlike atoms” and that in H-F the electrons are not shared equally. This causes The bond between H-F to have “ionic character” and this leads to a stronger bond that expected. Back to our Discussion on Hydrides Where are the electrons? c and bonding in binary hydrides - Classifiction d+ d- d+ d- Li H H F Li H H – H H F High electron Low electron density density  Lithium Hydride; an ionic / saline hydride Hydrogen Fluoride; a covalent hydride When hydrogen forms a binary compound to a When hydrogen forms a binary compound to a more electropositive element , like lithium (cp = more electronegative element, like fluorine (cP = 0.98). The hydrogen atom is described as 4.00). The hydrogen atom is described as protic. hydridic. Oxidation state of H = +1 Oxidation state of H = -1 High IE of H (+1312 kJmol-1) means its difficulty Favourable electron gain enthalpy = -73 kJmol-1 to produce an H+ ion. Binary hydrides of groups 1 and 2 are hydridic Protic hydrides are covalent molecules. and mostly ionic solids Electronegativity and the Reactions of Binary Hydrides Let's define a bond of any element to hydrogen as an E–H bond. If the electronegativity of E ≈ H, then the breaking of the E–H bond tends to be homolytic. Forming an H atom and a radical. For example, the combustion of hydrocarbons. E–H → E· + H· If E is more electronegative than H, heterolytic cleavage occurs, releasing a proton, protonic. The compound behaves as a Brønsted acid and can transfer H to a base. E–H → E- + H+ Heterolytic bond cleavage also occurs in compounds where E is less electronegative than H, including saline hydrides. In this case, the H atom is hydridic, and an H ion is transferred to a Lewis acid. E–H → E+ + H-  Ionic/ Saline Hydrides - Groups 1 and 2 Ø Salt-like, high-melting, white, crystalline stoichiometric hydrides with considerable ionic character. Ø LiH to CsH all adopt a rock salt (NaCl) structure. MgH2 adopts a rutile structure and CaH2 – BaH2 adopts a PbCl2 structure. Ø Formed by the group 1 and group 2 metals; Li – Cs and Ca—Ba 2Na (s) + H2 (g) 2NaH (s) Prepared by direct reaction of the elements at 300 – 700 oC Mg (s) + H2 (g) MgH2 (s) Ø They are mostly ionic compounds and conduct electricity when molten. Ø Ionic character increases down the group. Ø Reactivity increases down the group Saline Hydrides - Reactivity The hydride anion is relatively large (154 pm). The single charge of the hydrogen nucleus “can barely manage to keep control of the two electrons in the hydride anion, and they are easily lost. The H- anion is very polarizable and the combination of H- with a small polarizing cation such as Li+ and Mg2+ leads to a high degree of covalent character in the bond. The H- anion is a good proton acceptor (Bronsted- Lowry base) Ionic hydrides of group 1 are very reactive and often stored in oil Ionic hydrides react with water to form H2 and an alkaline solution. Good reducing agents: Eo (H2/H-) = -2.25 V Eo (Na+/ Na) = - 2.7 V Covalent Hydrides Covalent or molecular hydrides are binary hydrides formed with beryllium and p-block elements. The difference in bond polarity in these compounds results in a range of reaction types and hydrogen can be classified as H +, H -, or H. Hydrogen forms Strong Covalent Bonds. Twenty Strongest Covalent Bonds o Covalent bonds between hydrogen and other elements are among the strongest in chemistry. o Hydrogen ranks with fluorine and oxygen in its ability to form strong covalent bonds. o Of the twenty strongest covalent bonds, nine involve hydrogen. o Hydrogen forms stronger bonds with period two elements (B, C, N, O, F) than with period three elements in the same group (e.g., H-F vs H-Cl). o For polar bonds such as H–N, H–O, H–F, and H–Cl, bond strength can be partially attributed to the ionic nature of these bonds. This, however, does not explain the strong non-polar bonds that hydrogen forms. o Hydrogen’s ability to form strong bonds with non-polar compounds is attributed to its ability to create stable molecular orbitals involving the overlap of hydrogen's 1s orbital with most other atomic orbitals. Hydrogen forms Strong Covalent Bonds. Twenty Strongest Covalent Bonds Excellent orbital overlap involving hydrogen's 1s orbital results in energetically stable bonds. This explains the unreactive nature of H2. Hydrogen forms Strong Covalent Bonds. Twenty Strongest Covalent Bonds Weak E–H bonds formed with heavier p-block elements are due to poor overlap between the compact hydrogen 1s orbital and the diffuse s and p orbitals on E E = element Classification of Covalent Hydrides Most Common hydrides of the s- and p- block Electron-deficient hydride: not possible to describe the bonding of the structure in terms of 2-centre 2-electron bonds (e.g. BeH2 and group 13 hydrides. Electron –precise: all the valence electrons of the central atom are involved in forming bonds (grp 14) Electron-rich: Not all the electrons on the central atom participate in bonding. Groups 15-17 Types of Covalent Hydrides 1. Numerous covalent hydrides of boron, e.g. B2H6, B4H10 and B5H11 2. The three-dimensional polymeric covalent hydride of BeH2 3. Neutral, binary XH4 compounds of group 14. e.g. CH4 4. Somewhat basic binary XH3 compounds of group 15, e.g. NH3 and PH3 5. Weakly acidic or amphoteric binary XH2 compounds of group 16, e.g. H2O and H2S 6. Strongly acidic, binary HX compounds of group 17, e.g. HF and HI 7. Hydridic, complex compounds of hydrogen, e.g. LiAlH4 and NaBH4  Beryllium Hydride: BeH2 has a 3D covalent polymeric structure – Group 2 anomaly BeH2 is relatively stable in water, but the other Group 2 hydrides all react to give hydro- gen gas and alkaline solutions. For example, Colorless solid CaH 2 (s) + 2 H2 O (1) Ca(OH) 2 (aq) + 2 Hz (g) Bridging hydrogen between beryllium atoms – cp Calcium hydride is used in laboratories as a drying agent for organic solvents such as etha- nonitrile (acetonitrile, MeCN) and dichloromethane (CH 2 Cl 2). 3-centre 2-electron bond BeH2 is relatively stable in water – evidence of - Be - / covalency....._ Be / ""- H "- B / e " H Be2+ is a small cation with a high charge density I ,/ \ H H I / __ Be -- Be - \ this affects its ability to form a lattice. I "'. Be also has a high electronegativity value Figure 25.2 Solid BeH 2 forms a three-dimensional polymeric structure with bridging hydrogen atoms. compared to other s-block elements Beryllium hydride is an electron deficient hydride Be2+ is strongly polarizing and draws electron pair towards itself. H H 2 Electrons Centre 2 B H B H Centre 1 H H Two Centre Two Electron Bond, (2c,2e bond)  In a 3c-2e bond, the two electrons are 2 centre shared by three atoms with each atom contributing one orbital to the formation of the bond. Note the two different B-H bond 1 centre 2 electrons lengths in diborane. 3 centre Three Centre Two Electron Bond, (3c,2e bond) o Let's consider two sp2 hybridised BH3 fragments coming together. Each boron atom would use its sp2 hybrid orbitals and two electrons to form B - H bonds. o A B–H bond can “donate” two electrons into the empty p orbital on boron, forming a bridging hydrogen. And the formation of a three-centre two-electron bond. o A full arrowhead would indicate that hydrogen is donating 2 electrons using a dative covalent bond. A better descriptor used a half-arrow to suggest that the electrons come from the B–H bond. In B2H6 boron is sp3 hybridized Three Centre Two Electron Bond, (3c,2e bond) We can also conceptualize the bonding in diborane by viewing two sp3 hybridised BH2 fragments, each with a single electron in an sp3 orbital and the fourth hybrid orbital being empty. Two sp3 orbitals on boron can overlap with a hydrogen orbital to form two 2 centred, 2 electron B-H bonds. H B Forms two 2c,2e B bonds H 2 Hydrogen Boron - sp3 atoms - 1s1 hybridizaed with 3 electrons A bridging B-H bond is formed when the orbital of a third hydrogen atom overlaps with both an empty sp3 orbital from one BH2 fragment and with another singly occupied sp3 hybrid of another BH2 fragment. This results in the formation of a three-centre two- electron hydrogen bridge bond.  The 3-Centre 2-electron bond (3c-2e) – Bonding in diborane In a 3c-2e bond, the two electrons are shared by three atoms with each atom contributing one orbital to the formation of the bond. Note the two different B-H bond lengths in diborane. The simplest hydride of boron is BH3 (borane). Sp2 hybridised with an empty p-orbital. Only observed in the gas phase. BH3 readily dimerises to form B2H6. - How many electrons are present in diborane? How many electrons are required so that each element in diborane obeys the octet rule?  Group 13 Hydrides (EH3 or E2H6 ) Continued Electron deficient hydrides - Diborane is a gas at room temperature. - B2H6 combusts in air to produce boric acid: B2H6 (g) + 3 O2 (g) 2 H3BO3 (s) diborane - B2H6 hydrolyses in water to produce boronic acid and hydrogen B2H6 (g) + 6H2O (g) 2H3BO3 (s) + H2 (g) Decaborane (B10H14) - Boron forms many other boranes including decaborane which is produced when B2H6 is heated to 100oC. - Hydrides of other group 13 elements are not as stable as the boron hydrides. - AlH3 is a polymer with bridging hydrogens and decomposes above 150oC. AlH3 – Alane, polymer - Ga2H6 is similar in structure to B2H6, and is unstable at room Decomposes at 150 oC temp. Indium and thallium hydrides are only stable at low temp. Methane; CH4 (g) , DfH◦ = -74.8 kJmol-1 Group 14 hydrides (EH4) Ni cat. 400 C CO2 + 4H2 CH4 + 2H2O pressure Reactions Precise hydrides CH4 + 2O2 CO2 + H2O Carbon forms an unlimited about of hydrides i.e. hydrocarbons (Organic chemistry – Alkanes, alkenes, alkynes, etc.) Other group 14 elements form stable tetrahydrides – SiH4, GeH4, and SnH4. Except plumbane (PbH4) which is unstable. Silane; SiH4 (g), DfH◦= +34.3 kJmol-1 Like the hydrides of group 13, group 14 hydrides become less stable as you descend the group. SiCl4 + LiAlH4 SiH4 + AlCl3 + LiCl DfH◦ values of the non-carbon hydrides are positive indicating that these hydrides can not be prepared directly from their elements. Reactions spontaneous Silanes and germanes with chains up to 10 silicon or germanium atoms SiH4 + 2O2 SiO2 + H2O combusion have been isolated and characterized. More reactive than their carbon counterparts semiconductor SiH4 500 C Si + 2H2 grade silicon  Group 15 hydrides; EX3 – Electron Rich Polar covalent molecule, Nonpolar covalent molecule, c N = 3.04 c H= 2.20 c P = 2.19 c H= 2.20 PH3 and the other group 15 hydrides are poor Lewis bases. Bronsted- Lowery Base: Decrease in H-X-H bond angle indicates that lone pairs do not reside in hybridized sp3 orbitals but rather in an unhybridized s orbital and that p orbitals are used for P-H Electron pair donor – Lewis Base: bonds. S orbitals are not directional the lone pair is more diffused and is less available for donation to Lewis acids.  Water is a liquid at room temperature while all the other dihydrides are toxic gases. Group 16 High melting point and boiling point in water due to hydrogen bonding. Binary Similar to group 15 hydrides the bond angles of 16 Hydrides H2E hydrides decrease down the group.  CHEM1910 Block 1: Chemistry of the Main Group Elements Lecture 4 February 1st 2024 Dr Nickeisha Stephenson Covalent Hydrides Covalent or molecular hydrides are binary hydrides formed with beryllium and p-block elements. The difference in bond polarity in these compounds results in a range of reaction types and hydrogen can be classified as H +, H -, or H. Classification of Covalent Hydrides Most Common hydrides of the s- and p- block Electron-deficient hydride: not possible to describe the bonding of the structure in terms of 2-centre 2-electron bonds (e.g. BeH2 and group 13 hydrides. Electron –precise: all the valence electrons of the central atom are involved in forming bonds (grp 14) Electron-rich: Not all the electrons on the central atom participate in bonding. Groups 15-17 Types of Covalent Hydrides 1. Numerous covalent hydrides of boron, e.g. B2H6, B4H10 and B5H11 2. The three-dimensional polymeric covalent hydride of BeH2 3. Neutral, binary XH4 compounds of group 14. e.g. CH4 4. Somewhat basic binary XH3 compounds of group 15, e.g. NH3 and PH3 5. Weakly acidic or amphoteric binary XH2 compounds of group 16, e.g. H2O and H2S 6. Strongly acidic, binary HX compounds of group 17, e.g. HF and HI 7. Hydridic, complex compounds of hydrogen, e.g. LiAlH4 and NaBH4 Draw the Lewis structure of beryllium hydride. Does hydrogen obey the duet (octet) rule? H Be H Does Beryllium obey the octet rule? How would you classify BeH2? Can you sketch the structure of BeH2? Beryllium hydride is a covalent molecule and has a network structure How does this differ from the other binary hydrides of group 2? Beryllium Hydride Continued The Be2+ ion is highly polarizing. Due to its charge and small size: High charge density As a result, beryllium compounds tend to be covalent and their properties can be attributed to covalent bonding rather than ionic bonding. The chemistry of Be resembles the chemistry of aluminium more than the chemistry of the other group two elements: Diagonal relationship The strong polarizing power results in moderately covalent compounds, and its small size limits to four the number of groups that can attach to the ion. These two features together are responsible for the prominence of the tetrahedral BeX4 unit. Similar structure seen in BeCl2 Atkins-Chemical Principles 5th Ed.  Beryllium Hydride: BeH2 has a 3D covalent polymeric structure – Group 2 anomaly BeH2 is relatively stable in water, but the other Group 2 hydrides all react to give hydro- gen gas and alkaline solutions. For example, Colorless solid CaH 2 (s) + 2 H2 O (1) Ca(OH) 2 (aq) + 2 Hz (g) Bridging hydrogen between beryllium atoms – cp Calcium hydride is used in laboratories as a drying agent for organic solvents such as etha- nonitrile (acetonitrile, MeCN) and dichloromethane (CH 2 Cl 2). 3-centre 2-electron bond BeH2 is relatively stable in water – evidence of - Be - / covalency....._ Be / ""- H "- B / e " H Be2+ is a small cation with a high charge density I ,/ \ H H I / __ Be -- Be - \ this affects its ability to form a lattice. I "'. Be also has a high electronegativity value Figure 25.2 Solid BeH 2 forms a three-dimensional polymeric structure with bridging hydrogen atoms. compared to other s-block elements Beryllium hydride is an electron deficient hydride Be2+ is strongly polarizing and draws electron pair towards itself.  Draw the Structure of Boron Hydride Does hydrogen obey the duet (octet) rule? H Does Boron obey the octet rule? H B H How would you classify BH3? Can you sketch the structure of Boron Hydride? Diborane H H 2 Electrons Centre 2 B H B H Centre 1 H H Two Centre Two Electron Bond, (2c,2e bond)  In a 3c-2e bond, the two electrons are 2 centre shared by three atoms with each atom contributing one orbital to the formation of the bond. Note the two different B-H bond 1 centre 2 electrons lengths in diborane. 3 centre Three Centre Two Electron Bond, (3c,2e bond) o Let's consider two sp2 hybridised BH3 fragments coming together. Each boron atom would use its sp2 hybrid orbitals and two electrons to form B - H bonds. o A B–H bond can “donate” two electrons into the empty p orbital on boron, forming a bridging hydrogen. And the formation of a three-centre two-electron bond. o A full arrowhead would indicate that hydrogen is donating 2 electrons using a dative covalent bond. A better descriptor used a half-arrow to suggest that the electrons come from the B–H bond. In B2H6 boron is sp3 hybridized Three Centre Two Electron Bond, (3c,2e bond) We can also conceptualize the bonding in diborane by viewing two sp3 hybridised BH2 fragments, each with a single electron in an sp3 orbital and the fourth hybrid orbital being empty. Two sp3 orbitals on boron can overlap with a hydrogen orbital to form two 2 centred, 2 electron B-H bonds. H B Forms two 2c,2e B bonds H 2 Hydrogen Boron - sp3 atoms - 1s1 hybridizaed with 3 electrons A bridging B-H bond is formed when the orbital of a third hydrogen atom overlaps with both an empty sp3 orbital from one BH2 fragment and with another singly occupied sp3 hybrid of another BH2 fragment. This results in the formation of a three-centre two- electron hydrogen bridge bond.  The 3-Centre 2-electron bond (3c-2e) – Bonding in diborane In a 3c-2e bond, the two electrons are shared by three atoms with each atom contributing one orbital to the formation of the bond. Note the two different B-H bond lengths in diborane. The simplest hydride of boron is BH3 (borane). Sp2 hybridised with an empty p-orbital. Only observed in the gas phase. BH3 readily dimerises to form B2H6. - How many electrons are present in diborane? How many electrons are required so that each element in diborane obeys the octet rule?  Group 13 Hydrides (EH3 or E2H6 ) Continued Electron deficient hydrides - Diborane is a gas at room temperature. - B2H6 combusts in air to produce boric acid: B2H6 (g) + 3 O2 (g) 2 H3BO3 (s) diborane - B2H6 hydrolyses in water to produce boronic acid and hydrogen B2H6 (g) + 6H2O (g) 2H3BO3 (s) + H2 (g) Decaborane (B10H14) - Boron forms many other boranes including decaborane which is produced when B2H6 is heated to 100oC. - Hydrides of other group 13 elements are not as stable as the boron hydrides. - AlH3 is a polymer with bridging hydrogens and decomposes above 150oC. AlH3 – Alane, polymer - Ga2H6 is similar in structure to B2H6, and is unstable at room Decomposes at 150 oC temp. Indium and thallium hydrides are only stable at low temp. Methane; CH4 (g) , DfH◦ = -74.8 kJmol-1 Group 14 hydrides (EH4) Ni cat. 400 C CO2 + 4H2 CH4 + 2H2O pressure Reactions Precise hydrides CH4 + 2O2 CO2 + H2O Carbon forms an unlimited about of hydrides i.e. hydrocarbons (Organic chemistry – Alkanes, alkenes, alkynes, etc.) Other group 14 elements form stable tetrahydrides – SiH4, GeH4, and SnH4. Except plumbane (PbH4) which is unstable. Silane; SiH4 (g), DfH◦= +34.3 kJmol-1 Like the hydrides of group 13, group 14 hydrides become less stable as you descend the group. SiCl4 + LiAlH4 SiH4 + AlCl3 + LiCl DfH◦ values of the non-carbon hydrides are positive indicating that these hydrides can not be prepared directly from their elements. Reactions spontaneous Silanes and germanes with chains up to 10 silicon or germanium atoms SiH4 + 2O2 SiO2 + H2O combusion have been isolated and characterized. More reactive than their carbon counterparts semiconductor SiH4 500 C Si + 2H2 grade silicon  Group 15 hydrides; EX3 – Electron Rich Polar covalent molecule, Nonpolar covalent molecule, c N = 3.04 c H= 2.20 c P = 2.19 c H= 2.20 PH3 and the other group 15 hydrides are poor Lewis bases. Bronsted- Lowery Base: Decrease in H-X-H bond angle indicates that lone pairs do not reside in hybridized sp3 orbitals but rather in an unhybridized s orbital and that p orbitals are used for P-H Electron pair donor – Lewis Base: bonds. S orbitals are not directional the lone pair is more diffused and is less available for donation to Lewis acids.  Water is a liquid at room temperature while all the other dihydrides are toxic gases. Group 16 High melting point and boiling point in water due to hydrogen bonding. Binary Similar to group 15 hydrides the bond angles of 16 Hydrides H2E hydrides decrease down the group.  Group 15 hydrides; EH3 – Electron Rich Polar covalent molecule, Nonpolar covalent molecule, c N = 3.04 c H= 2.20 c P = 2.19 c H= 2.20 PH3 and the other group 15 hydrides are poor Lewis bases. Bronsted- Lowery Base: A decrease in the H-E-H bond angle indicates that lone pairs do not reside in hybridized sp3 orbitals but rather in an unhybridized s orbital and that p orbitals are used for P- Electron pair donor – Lewis Base: H bonds. S orbitals are not directional. The lone pair is more diffused and is less available for donation to Lewis acids.  Water is a liquid at room temperature, while all the other dihydrides are toxic gases. Group 16 High melting point and boiling point in water due to hydrogen bonding. Binary Like group 15 hydrides, the bond angles of 16 Hydrides H2E hydrides decrease down the group. Periodic Trends in the Acid Strength of group 16 and 17 Binary Hydrides Group 17 - Hydrohalic acids: a Closer Look Ø For hydrohalic acids, a decrease in bond strength results in a stronger acid Ø HF is the only weak acid in the family of hydrohalic acids. This is due to the unusual strength of HF covalent bond.  HX H-X bond length pKa D(H-X) / (pm) kJmol-1 HF 91.7 3.2 +570 HCl 127.5 -7 +431 HBr 141.5 -9 +366 HI 160.9 -10 +298 HX (aq) + H2O (l) X- (aq) + H3O+ (aq) 1) Decrease in HX bond dissociation energy makes the forward reaction more favourable as the group is descended. 2) The larger the X- anion, the weaker the attraction between X- and H3O+ ions. This means the reverse reaction becomes less favourable as the group is descended. 3) The large difference in acidity between HF and HCl is also due to very strong hydrogen bonds between F- and H3O+ ions which reduces the concentration of H3O+ ions in HF. Bond Enthalpies for X-H bonds in Binary Hydrides Trend in hydride stability is related to a reduction in E-H bond enthalpy As you descend a group, the atomic radii of E increase, and this leads to less efficient overlap between the E orbitals with hydrogen. Al and B have relatively weak bonds with hydrogen. What do these values tell us about the reactivity of these binary hydrides? Reactivity of E-H bonds: An Example Boron and aluminium form two fundamentally important hydrides that have numerous applications as reducing agents in organic and inorganic chemistry. These are sodium borohydride, NaBH4, and lithium aluminium hydride, LiAlH4. The anion in each case can be thought of as a hydride anion coordinated to a Lewis acid BH4- = H- + BH3 What does the E-H bond enthalpy tell us about the reactivity of the hydrides? B–H = 373 kJ mol-1 Al–H= 287 kJ mol-1 Sodium borohydride (NaBH4) and lithium aluminium hydride (LiAlH4) are used as reducing agents A covalent bond between hydrogen and B and Al. An ionic bond between the hydride complex and the alkali metal cation. o The Al-H bond is more polar than the B-H bond. o The Al-H bond is more hydridic than the B-H bond. o Looking at the relative bond lengths, we can say that the Al-H bond is weaker and longer than the B–H bond What does the E-H bond enthalpy tell us about the reactivity of the hydrides? Reaction with protic solvents: Protic hydrogen LiAlH4 is a stronger reducing agent than NaBH4. The Al-H bond is more polar than the B-H bond. The Al-H bond is more hydridic than the B-H bond. Looking at the relative bond lengths, we can say that the Al-H bond is weaker and longer than the B–H bond  CHEM1910 Transition Metals Lecture IV March 5th 2024 Dr Nickeisha Stephenson Lecture # Topic Reading* 1 (Feb 20th ) Metals of the d-block and pp 1256 - 1264 their properties 3 (Feb 27st ) Coordination Complexes: pp 1265 - 1272 geometries, oxidation state, and ligands. 4 (Mar 5th ) Introduction to Crystal Field Theory (CFT) – Colour, pp 1279 - 1295 Electronic spectra, and magnetism 5 (Mar 8th) CFT Continues pp 1279 - 1295 6 (Mar 12th ) Stability Constants pp 1295 - 1300 7 (Mar 14th ) Isomerism in coordination pp 1272 - 1277 complexes and naming * Chemistry3 by Burrows et al. Chapter 28  Square Planar Complexes – CN =4 and d8 Predict which of the following complexes could be square planar…. Colour in Transition metal Complexes When we discuss colour we are referring to the way in which matter interacts with frequencies of light in the visible spectra. Visible light falls in the range from about 380 – 750 nm Ø Before we can discuss the colour of transition metal complexes, we must first discuss how colour is perceived by the human eye. In the absence of a “filter” the Red light is reflected from the surface totality of the visible spectra is of the object, stimulating the cone Reflected light can be made up of perceived as white. cells in our eyes to send signals to the several wavelengths of light and not brain for decoding – red block just a single wavelength – brown table Ø When a substance absorbs certain wavelengths of light in the visible region of the electromagnetic spectrum, the colour of the substance is determined by the wavelength of light that remains. What Colour is the compound shown in the Vis light absorption spectrum shown below? What Colour is the compound shown in the Vis light absorption spectrum shown below? Answer: Orange Molecule is carotene which is the pigment found in carrots  Observed colour is the complementary colour of that absorbed..... it appears this colour Carotene absorbs photons mostly in the blue region of the spectrum and some what in the green region. maximum absorption (lmax) ~450 nm if a substance absorbs here....  Colour in Coordination Compounds Ø Unlike metals of the s and p block, colour is a characteristic feature of d-block metal complex ions. If 1% of Al3+ is replaced by Cr3+ , the chromium ion will absorb strongly in the blue-violet and yellow-green region of the visible spectrum. Red light is M+ transmitted giving us the characteristic colour of a ruby. M+ = Fe3+ results in yellow topaz, Corundum Corundum crystal structure M+ = Fe2+, Fe3+ and, Ti4+ results in Al2O3 Hard colourless mineral sapphires - blue when, M+ = Al Transition Metals are Coloured because of the Splitting of the d-orbitals Introduction to crystal field theory. In an octahedral complex, the d orbitals are split into two groups, called the eg set (dz2 and dx2-y2) and the t2g set (dxy, dyz, dxz). The gap between the orbitals is symbolized as D and can be used to explain the colour and magnetic properties of first-row transition metal complexes. Colour and the Spectrochemical Series https://www.youtube.com/watch?v=w3UgDC9nkf0  Ni(dmg)2 NC CN Ni NC CN What wavelength / colour of light is being absorbed by the above complexes?  Ø Colour is the result of electronic transitions occurring between molecular orbitals. Ø The colour observed in d-block complexes can be explained using crystal field theory. Ø For many transition metal complexes, D the crystal field splitting energy, corresponds to the energy of a visible light photon. Ø When white light passes through a solution containing [Ti(H2O)6 ]3+ ions, photons of yellow and green light are absorbed Ø The energy of the light absorbed corresponds to excitation of an electron from a t2g orbital to an eg orbital. [Ti(H2O6)]3+ A spectrophotometer can be used to observe the electronic transitions occurring when light of a certain wavelength interacts with matter  Colour in Transition Metal Complexes A given photon of light can be absorbed by a molecule only if the wavelength of light provides the energy required by the molecule to bring about a specific electronic transition. The wavelength absorbed is determined by: Energy DE = hn = hc / l Do = hc / l Where h = plank’s constant, c = the speed of The energy of this part of the visible light, and l is the maximum wavelength being spectrum is just right to promote the absorbed. electron from the t2g orbital to the eg Therefore, if you know lmax you can determine orbital the magnitude of Do The complex [Ti(OH2)6]3 absorbs light of wavelength 510. nm. What is the ligand field splitting in the complex in kilojoules per mole (kJ mol -1)? The complex [Ti(OH2)6]3 absorbs light of wavelength 510. nm. What is the ligand field splitting in the complex in kilojoules per mole (kJ mol -1)? A CLOSER LOOK AT THE METAL LIGAND INTERACTION □ Cobalt (III) – d6 metal. CN = 6 so this is an octahedral complex 3+ NH3 How many electrons do the d-orbitals hold? NH3 NH3 Co NH3 NH3 How many electrons will each ammonia ligand donate NH3 to cobalt? Where do these electrons go? 3d  Hybridization Theory Electrons from ligand occupy empty, hybrid metal orbitals.  Hybridization Theory Empty 4s and 4p will hybridize to form four new sp3 atomic orbitals Transition metal complexes can show magnetic properties. Magnet on: Diamagnetic Coordination complexes Gouy Balance with paired electrons are repelled by a magnetic field  OH2 CN H2O OH2 CN Fe SO4 NC H2O OH2 4K Fe CN OH2 NC CN Ø What is iron's oxidation state and d electron count in the coordination compounds shown? Ø How would you expect these Fe comppounds to behave on a Gouy Balance?  Hybridization Theory OH2 CN H2O OH2 CN Fe SO4 NC H2O OH2 4K Fe CN OH2 NC CN 3d sp3d2 4d Ans: Both complexes are Fe2+ complexes and based on the Valence Bond model we would expect them to be paramagnetic  LIMITATIONS OF THE LOCALIZED ELECTRON MODEL Valance bond theory can be used to rationalize ge