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Chapter 3

Chapter 3. Atoms: The Building Blocks of Matter. 3-1. The Atom: From Philosophical Idea to Scientific Theory. Foundations of Atomic Theory. Democritus - First to consider the smallest piece of matter. Around 400 BC he proposed the idea of the atom, which means indivisible. No proof.

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Chapter 3

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  1. Chapter 3 Atoms: The Building Blocks of Matter

  2. 3-1 The Atom: From Philosophical Idea to Scientific Theory

  3. Foundations of Atomic Theory • Democritus - First to consider the smallest piece of matter. • Around 400 BC he proposed the idea of the atom, which means indivisible. • No proof. • 1790’s scientist started using quantitative analysis of chemical reactions.

  4. The Law of Conservation of Mass • Mass is neither created or destroyed during ordinary chemical reactions. • The same elements are present in the same quantities before and after the reaction, they are just rearranged.

  5. Law of Definite Proportions • A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. • Composed of a fixed proportion of elements. • Ex: Sodium chloride, NaCl • 39.34% Na and 60.66% Chlorine by mass.

  6. Law of Multiple Proportions • If two or more different compounds are composed of the same 2 elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small, whole numbers. • Ex: Carbon monoxide, CO, 1g C & 1.33g O and Carbon dioxide, CO2, 1g C & 2.66g O. • The masses of the O is a 2 to 1 ratio.

  7. Dalton’s Atomic Theory • Proposed by John Dalton in 1808 as an explanation to the Laws of Cons. of Mass, Definite Proportions & Multiple Proportions. • He reasoned that elements were composed of atoms and only whole numbers of atoms can combine to form compounds.

  8. Dalton’s Atomic Theory • 1. All matter is composed of atoms. • Atoms of an element are identical in size, mass • and properties; atoms of different elements • differ in size, mass and properties. • 3. Atoms cannot be subdivided, created or destroyed

  9. Dalton’s Atomic Theory • Atoms of different elements combine in whole- • number ratios to form compounds. • In reactions, atoms are combined, separated or • rearranged.

  10. Dalton’s Explanation • Conservation of Mass • In the reaction C + O  CO, • the mass of CO is equal to the mass of C + • the mass of O and vice versa.

  11. Dalton’s Explanation • The Law of Definite Proportions • A given compound is always composed of the same combination of atoms.

  12. Dalton’s Explanation • Law of Multiple Proportions • C + O  CO and • C + O2  CO2 • The 2-1 ratio of oxygen masses results because carbon dioxide always contains twice as many atoms of oxygen as does carbon monoxide.

  13. Modern Atomic Theory • Atoms are divisible into smaller particles. • Elements can have atoms with different masses (isotopes). • Atomic theory has been revised. • Most important concepts: • All matter is composed of atoms. • Atoms of one element differ in properties from atoms of other elements.

  14. 3-2 The Structure of the Atom

  15. Remember • Atom – the smallest part of an element that retains the properties of that element. • Nucleus – very small center region of the atom that contains at least 1 proton (+) and usually 1 or more neutrons (neutral).

  16. Remember • Surrounding the nucleus is a region with negatively charged particles called electrons. • Protons, neutrons and electrons are called subatomic particles.

  17. Discovery of the electron • Late 1800’s • Experiments were preformed by running current through various gases at low pressures. • Were carried out in cathode ray tubes.

  18. When current was passed through the CRT, the surface of the tube directly opposite of the cathode glowed. • Hypothesis – the glow was caused by a stream of particles called cathode rays.

  19. Further Experimentation • Supported the existence of cathode rays. • Showed that cathode rays had enough mass to cause motion. • Cathode rays are negatively charged. • 1897 JJ Thomson concluded cathode rays are composed of negative particles, later named electrons.

  20. Charge & Mass of the Electron • Thomson’s experiment revealed electron’s have a very large charge compared to its tiny mass. • 1909 Robert Millikan determined the mass of an electron is about 1/2000 the mass of a simple H atom. More accurate; 9.109 X 10-31 kg, or 1/1837. • He also confirmed e- have a negative charge & all atoms have electrons.

  21. Assumptions made based upon the electron: • Because atoms are electrically neutral, they must contain positive charge to balance the electrons. • Because electrons have so much less mass than the atom, atoms must contain other particles that account for most of their mass.

  22. Discovery of the Atomic Nucleus • 1911 Ernest Rutherford with Hans Geiger & Ernest Marsden: • Bombarded gold foil with alpha particles. • Expected particles to pass through. • 1 in 8000 were redirected back to the source.

  23. Discovery of the Atomic Nucleus • Rutherford concluded that the force must be caused by a densely packed bundle of positive matter. • He called it the nucleus. • He also discovered that the volume of the nucleus is very small compared to the volume of the atom.

  24. Composition of the Nucleus • All atomic nuclei are made of 2 kinds of particles; protons & neutrons. • A proton has a positive charge equal in magnitude to the negative charge of an electron. • Atoms are neutrally charged because they contain equal numbers of protons & electrons.

  25. Forces in the Nucleus • Because like charges repel, the nucleus should be unstable. • There are strong attractions that occur when p-p’s or n-n’s are close. • Short-range prot-prot, prot-neut and neut –neut forces hold the nuclear particles together. • Called nuclear forces.

  26. 3-3 Counting Atoms

  27. General Info (Remember) • All atoms have the same basic parts. • Atoms of different elements have a different number of protons. • Atoms of the same element have the same number of protons.

  28. Atomic Number • Number of protons in the nucleus of each atom of an element. • Identifies the element • Remember that… • All atoms have the same basic parts. • Atoms of different elements have a different number of protons. • Atoms of the same element have the same number of protons.

  29. Isotopes • Atoms of the same element that have different masses. • Most have the same protons & electrons, but different neutrons. • Ex: All H atoms have 1 proton • Protium = 1 P & 1 E, 1 amu • Deuterium = 1 P & 1 N, 2 amu • Tritium = 1 P, 2 N & 1 E, 3 amu

  30. Isotopes of Hydrogen

  31. Determining Neutrons (Isotopes) • Mass # = Number of protons + neutrons. • Mass # - Atomic # = Neutrons • Ex: Chlorine-37 • 37 - 17 = 20 Neutrons • Nuclide – general term of any isotope of any element.

  32. Designating Isotopes • 2 Methods: • Hyphen Notation - Mass # written with a hyphen after the name. • Tritium is written as Hydrogen-3 • Nuclear Symbol – Chemical Symbol with • mass # as superscript & atomic # as subscript.

  33. Relative Atomic Mass • Standard used for atomic mass is the carbon-12 nuclide. • Assigned a mass of exactly 12 amus. • 1 amu = 1/12 the mass of a carbon-12 atom. • The atomic mass of any nuclide is determined by comparing it with the mass of a carbon-12 atom. • Isotopes have different masses, but don’t differ significantly in chemical behavior.

  34. Average Atomic Mass • Weighted aver of the atomic masses of the naturally occurring isotopes of an element. • Ex: 25% of marbles, 2.00g, 75% of marbles, 3.00g. • 25 X 2.00g = 50g • 75 X 3.00g = 225g • 275g / 100 • 2.75g

  35. Average Atomic Mass • or • Multiply the mass of each marble by the decimal representing its percentage in the mixture. • Add products together. • (2.00g X 0.25) + (3.00g X 0.75) = 2.75g

  36. Calculating Average Atomic Mass • Depends on the mass & relative abundance of each elements isotopes. • Ex: Copper • 69.17% Copper-63, 62.929 598 amu • 30.83% Copper-65, 64.927 793 amu • Multiply atomic mass of each isotope by the relative abundance in decimal form. • (0.6917 X 62.929 599 amu) + (0.3083 X 64.927 793 amu) = 63.55 amu

  37. Relating Mass to # of Atoms • The Mole (mol) • SI unit for amount of substance. • The amount of a substance that contains as many particles as there are atoms in exactly 12g of Carbon-12. • Counting unit (like dozens)

  38. Avogadro’s Number • Number of particles in exactly 1 mole of a pure substance. • Equals 6.022 X 1023 • Amedeo Avogadro • Alternative def for Mole – amount of a substance that contains avogadro’s number or particles.

  39. Molar Mass • The mass of 1 mole of a pure substance. • Measured in g/mol. • Equal to the atomic mass of the element in amus. • Ex: Li = 6.941 amu or 6.941 g/mol • Contains 1 mole of atoms • 6.941g of Lithium contains 1 mole of atoms.

  40. Gram to Mole Conversions Example: The molar mass of He is 4.00 g He/mole. How many grams of He are there in 2 moles? 2.00 mol He X 4.00 g He = 8.00 g He mol He

  41. Gram to Mole Conversions A chemist produced 11.9 g of aluminum. How many moles of aluminum were produced? 11.9 g Al X 1 mol Al = 11.9 mol Al = 0.441 mol Al 26.98 g Al 26.98

  42. Conversions with Avogadro’s # How many moles of silver are in 3.01 x 1023 atoms of silver? 3.01 X 1023 Ag atoms X mol Ag 6.022 X 1023 Ag atoms 3.01 X 1023mol Ag 0.500 mol Ag 6.022 X 1023

  43. Conversions with Avogadro’s # What is the mass in grams of 1.20 X 108 atoms of copper? 1.20 X 108 Cu atoms X mol Cu X 63.55 g Cu 6.022 x 1023 Cu atoms 1 mol Cu 76.23 X 1087.6 X 109 6.022 x 1023 1.2620 1.27 X 10-14

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