Chemistry 20A 2002 Review of key concepts from high school physics and chemistry Outline Physics Atomic building blocks subatomic particles Electrons protons neutrons Isotopes Relation between macroscopic and microscopic mole units Chemistry Molecular and Empirical formulas Stoichiometry Physics While chemistry may often seem like it is stamp collecting it is really governed by a few simple principles The most fundamental are mass conservation and energy conservation Mass conservation is useful in balancing atomic equations Energy conservation is very important since it tells us that energy goes somewhere For example in a reaction 2H2 O2 2H2O the potential energy in the reagents what you start from the left side of the equations is converted into a potential energy of the products water but some energy remains i e the potential energy of the products is lower than the potential energy of the reactants This extra energy is released in the form of kinetic energy i e some heat is released a more accurate statement would be formulated in 110A thermodynamics but this is good enough for our present purposes There is one extra principle which we would not tackle at this stage i e that entropy increases or in less fancy words that disorder increases Armed with these 3 principles it is possible to understand and sometimes to predict the outcome reagents remaining or products formed of much of chemical reactions and to understand the state that the products are in i e whether they are gas liquid solid etc The total energy is made of two parts Kinetic energy and potential energy E K E V Where K E 1 2 m v 2 while V is the potential energy The electrons are spread out over a region more or less spherical of size 1 3 1 1 Angstr m 10 10 m Quantum Mechanics later triumphed in predicting this size we can now see that it is true using special microscopes that can sample a single molecule So The spread out electrons determine size of atoms The electrons are very light compared to protons me mp and me 10 30 kg 1836 On a microscopic scale electromagnetic forces are much stronger than gravitational forces But gravity always has the same sign always attracts while electrostatic forces can cancel attract and repel For this reason gravity is stronger on a macroscopic scale Each of the subatomic particles has a different role in the chemistry The identity of an element is determined by Z which denotes the number of protons The number of neutrons may vary this leads to the phenomena of isotopes an element whose nuclei have the same number of protons but vary in the number of neutrons The effect of the neutrons is not felt beyond the nucleus that s why they don t directly influence the chemistry but they can be useful tools Examples Hydrogen Z 1 but Nn 0 1 2 p 1H 99 2H D Used in nuclear reactors 3H T Radioactive trace pn pn n Hydrogen bombs Carbon Z 6 Nn 6 7 8 12C mot common 13C important medically for MRI s 14C radioactive dating This is what you need to know now about subatomic particles Moles Macroscopic samples contain 1020 or more particles We therefore count these things atoms or molecules using the concept of a mole 6 1023 particles since we don t want to carry factors of 1020 in the equations Note the motivation for this rather arbitrary choice of 6 1023 and not say 1024 or 8 1023 1 gm of H atoms which exist as separate atoms and not as H2 molecules only at high temp contains 1 mole of atoms We define 6 1023 as Avogadro s Number and say that a mole contains Avogadro s number of things I e 1 gm of H atoms contains 1 mole of atoms Question how many moles does one gram of hydrogen contains at room pressure and temperature Answer only mole since at room pressure hydrogen is in H2 form Atomic and Molecular Masses Units macroscopically gram g and kilograms kg Don t confuse g with gravitational acceleration But in real life it is much easier to use 1 amu defined approximately as the mass of one proton or one hydrogen 1 amu 1g 1 7 10 24 g N Avog Now measuring masses is much more convenient m H2O 2 1 16 18 amu not exactly but almost g We can rewrite this as m H2O 18 mol 18 g mol 1 High school chemistry formulas and stoichiometry Now we ll review some simple aspects of high school chemistry which you do need to know Molecular formulas Tells the number of atoms in a molecule Example Methanol The molecular formula is CH4O This formula is not very useful since it does not indicate the structure In practice chemists would therefore often write the formula in a more transparent way for methanol it would be CH3OH With practice it is possible to get the structure for many molecules from this type of written formula Empirical Formulas and percentage composition 2 Empirical formula is the simplest formula for a compound but like molecular formula it lacks structure details Example hydrogen peroxide H2O2 molecular formula HO empirical formula For water H2O is both the empirical and molecular formula The empirical formula is very easy to determine just pick a sample of the compound and weight how many moles of each compound there are Example We find a gas sample containing 36g of C and 96g of O Determine percent C by mass percent C by moles Empirical formula Molecular formula Answer Percent C by mass 36 g 36 0 27 27 96 g 36 g 132 Percent C by moles First use MC 12 a m u and MO 16 a m u 36g n C 12 g mol 3 mol n O 96 g 6 mol 16 g mol The ratio is n C 0 5 1 2 n O 3 And Percent C by moles 3 6 0 33 Thus the empirical formula is CO2 Molecular formula cannot be determined on this basis Stoichiometry Balancing chemical equations Example C2H4 xO2 yCO2 zH2O We pick for convenience the coefficient of C2H4 to be 1 Determine stoichiometric coefficients x y z here y 2 since on left the coeff of C is 2 z 2 since on left the coeff of H is 4 So we are left with C2H4 xO2 2CO2 2H2O From which 2x 2 2 2 1 6 x 3 3