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CHEM Unit 4 Monday Oct 7 Greek philosophers Leucippus and Democritus Aristotle Skipping ahead – not on the test… Robert Boyle (about 1660) Jump to around 1800, give or take 20 years… Antoine Lavoisier Notes History of atomic theory about 400 BC Philosophers REASON, they don’t EXPERIMENT Two groups of philosophers o atomists o nonatomists 4 elements: earth, water, air, fire Leaders of the atomists If matter is divided enough times it reaches a point where it is made of indivisible particles – atoms Atoms of different elements have different sizes and shapes which make them behave differently Leader of the nonatomists Liked some things about the idea of atoms but had two big problems – unanswered questions: o What holds atoms together? o What’s in between atoms? Aristotle believed there could not be a location where there is nothing. Both sets of ideas were known and studied for the next 2000 years. HOWEVER, there weren’t many schools, not many people could even read, and most people who studied matter were ALCHEMISTS – studied matter more as if it was magic – they weren’t scientists. one of the first true scientists as opposed to alchemists (mostly) Wrote book – The Sceptical Chymist Two important things in this book o Defined element as matter that is made of fundamental, simplest particles that cannot be simplified any more (elements can’t be broken down by any chemical or physical change) o Said science was only valuable when every scientist shares not only his experimental results but the process he used to get them and his numerical results Farther of modern chemistry Wrote the first true chemistry textbook – Elements of Chemistry Identified the law of conservation of mass – mass can’t be created or destroyed; in any chemical or physical change, the mass you start with, you end with CHEM Unit 4 Notes History of atomic theory Lavoisier burns a diamond with a magnifying glass Joseph Proust Two ways to calculate % composition: Identified the law of constant composition (definite proportion) Any sample of the same compound will always contain the same elements in the same proportion by mass The percent by mass (mass percent) does not change. Mass percent can be calculated two ways: using the compound’s formula and the PT OR measuring the masses of elements in the compound in lab. You get the same answers. Use data collected in an experiment. o From ex. above: % comp. of MgO 𝑚𝑎𝑠𝑠 𝑀𝑔 o %𝑀𝑔 = 𝑚𝑎𝑠𝑠 𝑀𝑔𝑂 × 100 o %𝑀𝑔 = o %𝑂 = o %𝑂 = 5.00 𝑔 8.29 𝑔 𝑚𝑎𝑠𝑠 𝑂 × 100 = 60.3% 𝑀𝑔 𝑚𝑎𝑠𝑠 𝑀𝑔𝑂 3.29 𝑔 8.29 𝑔 × 100 × 100 = 39.7 % 𝑂 Use masses from PT o MgO contains 1 atom Mg and 1 atom O per molecule o Mass of 1 atom Mg = 24.305 u o Mass of 1 atom O = 15.9994 u o Mass of 1 molecule MgO = mass 1 atom Mg + 1 atom O = 24.305 u + 15.9994 u = 40.304 u 𝑚𝑎𝑠𝑠 𝑀𝑔 o %𝑀𝑔 = 𝑚𝑎𝑠𝑠 𝑀𝑔𝑂 × 100 o %𝑀𝑔 = o %𝑂 = o %𝑂 = 24.305 𝑢 40.304 𝑢 𝑚𝑎𝑠𝑠 𝑂 𝑚𝑎𝑠𝑠 𝑀𝑔𝑂 15.9994 𝑢 40.304 𝑢 × 100 = 60.3% 𝑀𝑔 × 100 × 100 = 39.7 % 𝑂 CHEM Unit 4 John Dalton Math of law of multiple proportions (you don’t have to DO it, but you do need to understand it) Dalton’s atomic theory (*essay question on test) Notes History of atomic theory Identified the law of multiple proportions If you have two different compounds that contain exactly the same elements, AND you take a sample of each element so that both elements contain exactly the same mass of element A, and then you look at the ratio of the mass of element B in compound 1 to the mass of element B in compound 2, the ratio is ALWAYS a simple wholenumber ratio. First person to measure and publish a table of atomic weights – got the idea from his law of multiple proportions. Developed the first complete atomic theory – ideas about atoms that EXPLAINED how matter behaved. Copper and oxygen form two compounds: copper (I) oxide, Cu2O copper (II) oxide, CuO If you take the amount of each compound that contains 10.00 g copper: The sample of Cu2O has a mass of 11.26 g. 10.00 g of this is copper; 1.26 g is oxygen. The sample of CuO has a mass of 12.52 g. 10.00 g of this is copper; 2.52 g is oxygen. Compare the masses of oxygen in each sample as a ratio: 1.26 g O : 2.52 g O Simplify the ratio by dividing both numbers by the smaller one: 1.26 𝑔 𝑂 2.52 𝑔 𝑂 ∶ 1.26 𝑔 𝑂 1.26 𝑔 𝑂 1∶2 The ratio is a simple whole-number ratio. All matter is made of tiny invisible indivisible particles called ATOMS. (definition of element – Boyle) All atoms of the same element are identical and different from atoms of every other element. (definition of element – Boyle; idea of atomic weight – Dalton) Atoms combine in simple whole-number ratios to form compounds. (law of constant composition – Proust) If atoms of the same element can combine in more than one ratio, they form more than one compound. (law of multiple proportions – Dalton) During chemical reactions, atoms break attachments, rearrange and form new attachments to different atoms but they are not created, destroyed or transformed into atoms of a different element. (law of conservation of mass – Lavoisier) CHEM Unit 4 Dalton’s model of the atom October 8, 2013 You-tube videos: Brian Cox, J. J. Thomson Notes History of atomic theory Billiard ball model Atoms are solid spheres Atoms of different elements have different masses. Atoms have little hooks on them that let them attach to other atoms. Cathode ray tube Beam of particles he shot went between two charged plates. The beam bent toward positive plate and away from negative plate. Opposite charges attract so particles in beam are negative Atoms are neutral Neutral atom gives off (contains) negative particles – very, very small Called them electrons Atom left behind became positive after electrons left called plum pudding model (like blueberry muffin) Blueberries are electrons – negative Bread is rest of atom - positive measured exact mass and charge of an electron gold foil experiment Shot alpha particles (small, positive particles) at gold foil Most went straight through A few bent sideways a few bounced backward WHY? Straight through atom is mostly empty Bent sideways pushed away by something positive (like charges repel) Thomson’s model Robert Millikan Ernest Rutherford CHEM Unit 4 Notes History of atomic theory Rutherford’s model – nuclear model Henry Moseley Protons James Chadwick Bounced back – hit something with a lot of mass and positively charged and bounced off Atom is mostly empty space Center is small, positive and massive – very dense – called nucleus Atoms are outside in large, mostly empty region called electron cloud. Discovered the atomic number – amount of positive charge in the nucleus. Predicted the atomic number matched a number of positive particles but died prior to proving it experimentally (it was later proved by Rutherford) Wrote periodic law – the properties of elements vary in a predictable repeating way based on atomic number Idea proposed by Moseley Experiments proving done by Rutherford due to Moseley’s death in WWI Discovered the neutron (how is really complicated, so don’t worry about it)