A Powerpoint on the use of hydrogen as a fuel reacting with Oxygen to generate electricity.

1. Chapter 10
Hydrogen

2. History
• Discovered about 200 years ago (1766)
Henry Cavendish (1731-1810)
Antoine Lavoisier (1743-1794)

3. History
• In 1931, hydrogen was discovered to have
isotopes
vs.
Harold C. Urey (1893-1981)
Nobel Laureate 1934
Frederick Soddy (1877-1956)
Nobel Laureate 1921

4. Isotopes of Hydrogen
• Three common isotopes:
– Protium (H)
• common hydrogen
• 99.985% abundant
– Deuterium (D)
• one neutron
• 0.015% abundant
– Tritium (T)
• two neutrons
• 1x10-15
% abundant

5. Physical Properties
• Vastly different among the three isotopes

6. Stability
• All are stable except for tritium
– synthesized in the universe and man-made by
neutron bombardment
– undergoes radioactive β decay

7. Nuclear Magnetic Resonance
• Studies nuclear spin

8. Nuclear Magnetic Resonance
• Peaks depend upon the chemical
environment

9. Nuclear Magnetic Resonance
• Intensity of the absorption depends upon
the nucleus
Nucleus Natural Abundance Relative Sensitivity
1
H 99.985 1.0
13
C 1.108 0.016
19
F 100 0.83
31
P 100 0.07

10. Magnetic Resonance Imaging
• Measures the 1
H nucleus in water

11. Properties of Hydrogen
• Periodic table placement
– reasons to place the element in both Group
1(alkali metals) and Group 17 (halogens)
Argument for Placement Argument Against
Placement
Alkali Metal Group forms monopositive ions is not a metal
has a single s electron does not react with water
Halogen Group is a nonmetal rarely forms a
mononegative ion
forms a diatomic
molecule
is comparatively
unreactive

12. Properties of Hydrogen
• Colorless, odorless gas
– m.p. -259°C
– b.p. -253°C
• Relatively non-reactive
– diatomic bond energy of 436 kJ/mol

13. Reactions of Hydrogen with
Diatomics
2H2(g) + O2(g) → 2H2O(g) (very fast)
H2(g) + F2(g) → 2HF(g) (very fast)
3H2(g) + N2(g) → 2NH3(g) (very slow)

14. Reduction Reactions of
Hydrogen
• Acts to reduce many metallic elements
CuO(s) + H2(g) → Cu(s) + H2O(g)
• Can also reduce double and triple bonds
with a catalyst
H2C=CH2(g) + H2(g) → H3C—CH3(g)

15. Preparation of Dihydrogen
• Reaction of dilute acids on metals
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

16. Preparation of Dihydrogen
• Steam reformer process
CH4(g) + H2O(g) → CO(g) + 3H2(g)
CO(g) + H2O(g) → CO2(g) + H2(g)
K2CO3(aq) + CO2(g) + H2O(l) → 2KHCO3(aq)

17. Hydrides
• Binary compounds of hydrogen
– has an intermediate electronegativity
• ionic hydrides
– LiH
• covalent hydrides
– HF
• metallic hydrides
– NiH2

18. Ionic Hydrides
• white solids
• metal cation and hydride ion
• very reactive
LiH(s) + H2O(l) → LiOH(aq) + H2(g)
• reducing agents
CaH2(s) + H2O(l) → Ca(OH)2(s) + H2(g)

19. Covalent Hydrides
• covalently bonds with all nonmetals and
weakly electropositive metals
• gases at room temperature
– hydrogen can be:
• nearly neutral
• substantially positive
• slightly negative

20. Neutral Covalent Hydrides
• low polarity
– only dispersion forces
• Examples:
– PH3
– CH4
– Hexene
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)

21. Positive Covalent Hydrides
• high melting and boiling points
– protonic bridging
• Examples:
– ammonia
– water
– hydrogen fluoride

22. Negative Covalent Hydrides
• Contains hydridic hydrogens
• Very reactive towards oxygen
• Examples:
– B2H6
– SiH4
– GeH4
GeH4(g) + 2O2(g) → GeO2(s) + 2H2O(l)

23. Borane Bonding
• possess bridging hydrogens
– hydridic bridges
– very reactive
B2H6

24. Borane Bonding
• Three-center, two-electron bond
– banana bonds

25. Borane Bonding
• MO Picture
+
B1 B2
σ
σ∗
σ
σ∗
σNB
H

26. Borane Bonding
• Other borane structures

27. Metallic Hydrides
• Hydrogen occupies the interstitial spaces
– non-stoichiometric
• TiH1.9
– less densities
– brittle
– lower conductivity

28. Synthesis and Reactions of
Metallic Hydrides
Ti(s) + H2(g) + heat/pressure → TiH1.9(s)
TiH1.9(s) + heat → Ti(s) + H2(g)
• applications in hydrogen storage

29. Uses of Metallic Hydrides
• Batteries
Cathode: Ni(OH)2(s) + OH-
(aq) → NiO(OH)(s) + H2O(l) + e-
Anode: [Ni-alloy](s) + H2O(l) + e-
→ [Ni-alloy]H(s) + OH-
(aq)

30. Water and Hydrogen Bonding
• Without hydrogen bonding, water would
melt at -100°C and boil at -90°C
• Liquid is denser than the solid

31. Water’s Phase Diagram
• Different from a normal phase diagram
Normal Water

32. Electrical Conductivity in Water
• highest for solutions of H3O+
and OH-

33. Clathrates
• a substance which is trapped in the crystal
lattice of another substance
– from “clathratus”
• enclosed behind bars
• gas hydrates
– methane
– noble gases

34. Biological Aspects of Hydrogen
Bonding
• Hydrogen’s properties play two key roles to
the existence of life
– closeness in electronegativity to carbon
– ability to form hydrogen bonds

35. Reaction Flowchart
• Shows the different types of reactions of a
certain species
H2
NaH
NH3
Cu
H2OHF
TiH1.9
O2F2
Na
CuO
N2Ti