Atomic Theory: Elements, Compounds and Stoichiometry Chemistry for Earth Scientists DM Sherman University of Bristol Compounds and Chemical Reactions! By the 18 th Century, it was realized that many substances were compounds formed by the combination of two or more elements. Etienne Francois Geoffroy 1672-1713 Page 1
The Elements as of 1808! The Elements as of 1937! Page 2
18 th Century Empirical Observations I! The Law of conservation of Mass: There is no detectable change in mass in an ordinary chemical reaction. + = 1.000 g Iron 0.432 g Oxygen 1.432 g Hematite 18 th Century Empirical Observations II! The Law of Constant Composition: A compound always contains the same elements in the same proportions by mass. = + 1.000 g Hematite 0.301 g Oxygen 0.698 g Iron Note: for some compounds/minerals, this is not true! Page 3
18 th Century Empirical Observations III! The Law of Multiple Proportions: The masses of one element that combine with a fixed mass of the second element are in a ratio of small whole numbers. Mineral Formula Mass of Iron (g) Mass of Oxygen (g) Wusite FeO 1.000 0.286 g Magnetite Fe 3 O 4 1.000 0.382 Hematite Fe 2 O 3 1.000 0.430 4/3 9/8 3/2 Dalton s Atomic Theory (1808)! Dalton explained these observations using atomic theory: Elements consist of tiny indivisible particles called atoms. Atoms of the same element are alike in mass and size. Atoms combine to form compounds in simple numerical ratios, such as 1:2, 2:3, etc. Atoms of two elements may combine in different ratios to form more than one compound. Page 4
Why is this so important?! Once we recognize that chemical compounds are made up of atoms, we can now express reactions that form compounds in terms of this idea: Fe + O = FeO (wustite) 2Fe+ 3O = Fe 2 O 3 (hematite) 3Fe + 4O = Fe 3 O 4 (magnetite) This makes everything much, much simpler! We can then predict the masses of things produced/ consumed by a chemical reaction if we know the relative atomic masses. Law of Combining Volumes! In 1808, Gay-Lussac proposed that gases, under equal conditions of temperature and pressure, react with one another in volume ratios of small whole numbers. Amadeo Avogadro (1811) hypothesized that equal volumes of gasses at the same P,T contain equal numbers of molecules (even if they are different elements!). Page 5
The Mole! Dalton had established a system of relative atomic masses based on hydrogen (relative mass = 1). Using Avagadro s hypothesis, Stanislao Cannizzaro (1861) was able to produce a system of relative atomic masses by comparing the relative masses of equal volumes of gas. The original definition of a mole is the number of oxygen atoms in 16.000 grams of oxygen. Relative Atomic Mass! With the invention of mass spectrometry (1912), it was discovered that atoms of a particular element may have several different masses! These are the isotopes. 78 Kr 0.356 % 80 Kr 2.27 % 82 Kr 11.6 % 83 Kr 11.5 % 84 Kr 57.0 % 86 Kr 17.3 % Page 6
Relative Atomic Mass! By convention, we take the isotope 12 C to be the standard: an atom of 12 C has a mass of 12 atomic mass units. The masses of other isotopes are given relative to this. The masses of the elements are the isotopic masses weighted by their isotopic abundances. e.g., Iron = 5.80% 54 Fe + 91.72% 56 Fe + 2.20% 57 Fe +.28% 58 Fe This gives a relative atomic mass of 55.85 g/mole. Masses of Elementary Particles! As we will learn, atoms consist nuclei (made up of protons and neutrons) surrounded by electrons. The isotopes of a particular element differ in the number of neutrons. Isotopic masses still don t come out to be exact integers because the masses of the protons and neutrons are not exactly equal: proton mass = 1.67262158 10-27 kg electron mass = 9.10938188 10-31 kg neutron mass = 1.67492729(28) 10 27 kg Page 7
Determination of Avogadro Constant! Faraday (1834) had determined the charge of a mole of electrons (F = 96485. Coulombs/mole) from electrochemical reactions. Mulliken (1910) had measured the charge of the electron e =1.602 10 19 C. Given the electron charge, we can now work out how many particles are in a mole: N A = F/e = 6.022 x 10 23 /mole We call this Avogadro s number. Molar Mass and Relative Molar Mass! Relative molar mass is a dimensionless quantity: RMM = Mass of 1 mole/(mass of 1 mole of 12 C) Molar mass is the mass (in grams) of 1 mole of something. Page 8
Moles/Mass and Chemical Calculations! Moles of A Moles of B Mass of A? Mass of B The Unit Factor Method! This is how you do chemical calculations. It can also be applied to many other problems. When you do any calculation: Always write out the units for each quantity. Treat the units like algebraic quantities. Make sure units cancel to give the desired units for the answer. This is, by far, the most important skill I can teach you!! Page 9
Moles/Mass and Chemical Calculations! Example 1: Consider the reaction: CH 2 O + O 2 -> CO 2 + H 2 O How many grams of CO 2 will be produced from 1 gram of CH 2 O? Solution: ( 1g CH 2 O) 1mole CH 2O 1 mole CO 2 44 g CO 2 30 g CH 2 O 1 mole CH 2 O 1 mole CO 2 = 1.47gCO 2 Moles/Mass and Chemical Calculations! Example 2: How many grams of Cu will be produced by smelting 1 g of chalcopyrite (CuFeS 2 ). The atomic mass of Cu is 63.5 g/mole and the molecular mass of CuFeS 2 is 183.6 g/mole. Solution: ( 1g CuFeS 2 ) 1mole CuFeS 2 1 mole Cu 63.5 g Cu 183.6 g CuFeS 2 1mole CuFeS 2 1 mole Cu = 0.346gCu Page 10
Important Concepts/Things to Learn! The elements Conservation of Mass The mole Elements and Isotopes Unit Factor Method Page 11