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Atoms – Building Blocks of Matter Notes - Chapter 3

Atoms – Building Blocks of Matter Notes - Chapter 3. FIRST CHEMICAL REACTION. I. The Atom: From Idea to Theory. A. 400 BC Democritus VS Aristotle

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Atoms – Building Blocks of Matter Notes - Chapter 3

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  1. Atoms – Building Blocks of Matter Notes - Chapter 3

  2. FIRST CHEMICAL REACTION

  3. I. The Atom: From Idea to Theory A. 400 BC Democritus VS Aristotle Democritus, an ancient Greek and student of Aristotle, proposed the 1st atomic theory he said that the world is composed of 2 things: void (empty space) and matter. No one supported him and he had NO experimental evidence to support his idea.

  4. Aristotleproposed that matter was composed of one continually flowing substance called hyle. This idea was widely supported and accepted until the late 1700’s and he too had NO experimental evidence to support his idea.

  5. What Accounts for Matter’s Secrets? The Greeks came up with at least 3 different theories ARISTOTLE – 4 ELEMENTS: AIR, EARTH, FIRE, WATER

  6. ALCHEMISTS ACCOMPLISHMENTS – NOT GOLD

  7. B. Late 1700’s Isaac Newton and Robert Boyle • It was not until the late 1700’s that anyone dared to question Aristotle’s wisdom. They suggested that Aristotle was incorrect but did not have their own theory to submit. • At this time chemist did believe, based on experiments, that there were different elements and that an element was a substance that could not be broken down by chemical means. Chemist knew that some substances could transform into different or new substances, they called this a chemical reaction.

  8. C. 1790’s - Basic laws that were established: Chemist also new, via improved balances, that when a chemical reaction occurred in a closed space that the mass of the material before the change equaled the mass of the material after the change. Now known as the Law of Conservation of Mass.

  9. Another realization was that substances always contained their elements in the same proportions by mass. For example: for any sample of sodium chloride, the mass of the sample is always 39.34% Na and 60.66% Cl. Now known as the Law of Definite Proportions.

  10. It was also known that elements combined to form more than one compound. Example: carbon monoxide and carbon dioxide. This is the Law of Multiple Proportions.

  11. D. 1803 John Dalton • British chemist who was the first to have a theory about matter being composed of atoms and how • atoms might look and behave. Dalton proposed an explanation for the Law of Conservation of Mass, Law of Definite Proportions, and Law of Multiple Proportions. He reasoned that elements were composed of atoms and that only whole numbers of atoms can combine to form compounds. He conceived on the atom as a solid billiard ball. Here is a summary of his theory:

  12. JOHN DALTON (1766 - 1844) REVIVES ATOMIC THEORY OF MATTER

  13. 1. All matter is composed of atoms. • 2. Atoms of the same elements are exactly the same and atoms of different elements are different. • 3. Atoms cannot be created, destroyed, or subdivided. • 4. Atoms of different elements combine in whole number ratios to form compounds. • 5. In chemical reactions, atoms are combined, separated, or rearranged.

  14. Democritus’s idea, because Dalton was able to relate atoms to the measurable property of mass, turned into a scientific theory!! The only aspect of Daltons’ Theory that is now known to be incorrect is the fact that atoms can be subdivided (into p+, e-, n). And that atoms of the same element can have deterrent masses (these are called isotopes).

  15. II. The Structure of Atoms Atom – smallest particle of an element that retains the chemical properties of that element. All atoms consist of 2 regions – the nucleus (p+ & n) and surrounding the nucleus is the electron cloud – a region occupied by the negatively charged particles called electrons. How do we know this?!

  16. 1. Discovery of Electron 1897 (by J.J. Thomson and Robert Millikan) 1st subatomic particle to be discovered – Thompson was working with electricity and magnetic fields. He was taking various gases and sending an electric current through the gas. When he did this he noticed that a glow was emitted. (What he was doing, he believed, was separating the electron from the nucleus of the gas atoms – this caused the glow!) Thompson went on to prove that the glow was actually a stream of negatively charged particles – called electrons. Symbol e-, charge –1, and mass of 0.00055amu (atomic mass unit, 1amu = 1.66 X10 –27 kg)

  17. Plum Pudding Model – Thompson proposed that the atom had negative electrons scattered throughout a positively charged area (proton area).

  18. 2. Protons 1919 (discovered by Rutherford/J.J. Thompson) Both Rutherford and Thompson knew that positively charged particles (protons) must exist (because an atom is neutral, if there is a negative charged electron then there has to be a positively charged proton to make it neutral.) They worked together to prove they existed. Proton symbol: +p, charge +1, mass 1.008 amu.

  19. 3. Discovery of the Atomic Nucleus 1911 Subatomic ParticleMassElectron 0.00055amuProton1.008 amu.Neutron1.008 amu. Discovered by Rutherford during his famous gold-foil experiment and realized that the main part of the atom’s mass is in the nucleus, and that it is positively charged. Summary of his experiment: • -Bombarded a thin piece of gold foil with positive alpha particles • -Most went through as though nothing was there • -Few (1 in 8000) ricochet back toward the source • -Few were deflected off to the side

  20. Rutherford’s Conclusion: the positive alpha particles had to have hit something else that was positively charged to cause the ricochet effect. The “something” was very small and dense because only a few hit it, therefore the atom must have a small positively charged nucleus, surrounded by mostly empty space (because most particles went through the gold foil.) New model of atom: • Electron 0.00055amu • Proton 1.008 amu. • Neutron 1.008 amu.

  21. 4. Neutrons 1932 (Proved by Chadwick) New something else existed in an atom because of the mass of the atom. Neutron is an electrically neutral particle, symbol n, mass equal that of protons.

  22. 5. Nuclear Forces – the +p and n stay close to each other due to these short-range forces that hold the +p and n together. Current Model of Atom:

  23. III. Counting Atoms Reading the periodic table 11 atomic number Na symbol 22.990 average atomic mass (in amu’s) Sodium name of element 23 mass number (the average atomic mass rounded to the nearest whole number)

  24. 1. Atomic Number • the number of protons in the nucleus. The atomic number identifies the element!!!!!!!!! Because atoms are neutral they contain the same number of electrons as protons. (Therefore the atomic number is the number of electrons as well.)

  25. 2. Atomic Mass – mass of 1 atom (measured in amu’s)

  26. 3. Mass Number • – the average atomic mass rounded to the nearest whole number, therefore it is the total number of protons and neutrons in an atom’s nucleus.

  27. Practice: How many protons are in each of the following? neutrons? electrons?

  28. 4. Isotopes • – atoms with the same number of protons (atomic number is the same) but different numbers of neutrons (mass number is different). Usually isotopes are referred to by their name (of symbol) and their mass number. Every element on the chart has at least 2 isotopes and some elements have as many as 25 isotopes. • Example: The isotopes of hydrogen have separate names rather than being called hydrogen-1, hydrogen-2, etc. Their names are protium (H-1), deuterium (H-2), and tritium (H-3). Diagram of protium, deuterium, and tritium:

  29. Practice: carbon-14, carbon-13, carbon-12

  30. Most elements occur naturally as mixtures of isotopes, as indicated in Table 3-4 of textbook. The percentage of each isotope in the naturally occurring element on Earth is nearly always the same, no matter where the element is found. The percentage at which each of an element’s isotopes occurs in nature is taken into account when calculating the element’s average atomic mass (which appears on the periodic table).

  31. Nuclide -a general term for any isotope of any element

  32. 5. Relative atomic masses • a.m.u.- atomic mass unit; One amu is exactly 1/12 the mass of a carbon-12 atom. So the atomic mass of any nuclide is determined by comparing it with the mass of the carbon-12 atom. The hydrogen-1 atom has an atomic mass of about 1/12 that of the carbon-12 atom, or 1 amu. 1 amu = 1.66X10-27kg

  33. 6. Average atomic mass • It is the weighted average of the masses of all the isotopes of that element. A weighted average reflects both the mass and the abundance of the isotopes as they occur in nature. Ex:ample: isotope Atomic mass abundance (%) • protium 1 99.985 • deuterium 2 0.015 • tritium 3 negligible • The average atomic mass of hydrogen is 1.0079. Multiply each mass number by the percent abundance and add them up.

  34. Practice: Element Z has 2 natural isotopes. The isotope with a mass number of 15 has a relative abundance of 30%. The isotope with a mass number of 16 has a relative abundance of 70%. Estimate the average atomic mass for this element.

  35. IV. Relating Mass to Numbers of Atoms 1. The Mole (can be abbreviated mol, but NOT m, which is the abbreviation for meter!) - the SI unit for amount of substance. A mole is the amount of a substance that contains as many particles as there are atoms in exactly 12 grams of carbon-12.

  36. 2. Avogadro’s Number -the number of particles in exactly one mole of a pure substance. This number was determined experimentally and its value is 6.022 X 1023, which means that 12 g of carbon-12 contains 6.022 x 1023 carbon-12 atoms.

  37. 3. Using the Mole and Avogadro’s Number • A mole can be thought of as a counting unit just like a dozen (12), gross (144), pair (2), ream (500), mole (6.022X1023 ).

  38. A. How many is a mole? • If every person living on Earth (6 billion people) worked to count out one mole of oranges (or anything else), and if each person counted continually at a rate of one orange per second, it would take about 4 million years for all the oranges to be counted! • If we had a mole of sand it would cover the earth 7 times over! If you had a mole of dollar bills, you could spend a million dollars every minute of your life and never spend it all!

  39. Since the mole is so large, we use it to count very tiny things – like atoms. Because the mole is so large, (and we now know that we cannot count out a mole of anything), how do we know when we have a mole of anything? • We determine the mass and relate that to the number of atoms present.

  40. 4. Molar Mass • – The mass of one mole of a pure substance. The pure substance can be an element or a compound. • The atomic mass is the mass of 1 atom of that element measured in amu’s. • The atomic mass is also equal to 1 mole of atoms measured in grams it is called the molar mass!!!! What a coincidence!!!!

  41. Mass of 1 atom of Pb = 207.2 amu • Mass of 1 mole of Pb atoms = 207.2 g • Mass of 1 atom of N = 14.01 amu • Mass of 1 mole of N atoms = 14.01 g • Mass of 1 atom of Ba = 137.33 amu • Mass of 1 mole of Ba atoms = 137.33 g • Mass of 1 atom of Al = 26.98 amu • Mass of 1 mole of Al atoms = 26.98 g

  42. Let’s prove it: Determine the mass, in grams, of 6.022X1023 atoms of aluminum. Use 1amu = 1.66X10-27kg.

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