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Understanding Atoms, Molecules, and Ions: Key Discoveries and Modern Atomic Structure

This chapter explores the fundamental concepts of atoms, molecules, and ions, detailing the atomic theory of matter developed by scientists like John Dalton and J.J. Thomson. It covers the laws of conservation of mass, constant composition, and multiple proportions. The discovery of electrons, protons, and neutrons, as well as radioactivity and the nuclear atom, are discussed along with their implications in modern atomic structure. The chapter emphasizes how atoms combine to form compounds and the foundational principles governing chemical reactions.

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Understanding Atoms, Molecules, and Ions: Key Discoveries and Modern Atomic Structure

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  1. CHAPTER 2 • Atoms, Molecules and Ions

  2. Atoms • Each element composed of atoms • All atoms of a given element are identical • Atoms of an element are not changed during a chemical reaction (the atoms just move from one chemical species to another) • Compounds are formed when atoms of more than 1 element combine

  3. Atoms • John Dalton • The Atomic Theory of Matter • Credited with developing Natural Laws

  4. Atoms • Law of conservation of mass • Law of constant composition • Law of multiple proportions: When two elements form different compounds, the mass ratio of the elements in one compound is related to the mass ratio in the other by a small whole number.

  5. Discovery and Properties of Electrons • Humphrey Davy (early 1800’s) - passed electricity through compounds • compounds decomposed into elements • compounds are held together by electrical forces • Michael Faraday - (1832-1833) - amount of reaction that occurs during electrolysis is proportional to current passed through compounds • Matter (atoms) is electrical in nature.

  6. Discovery and Properties of Electrons • Cathode Ray Tubes - (late 1800’s & early 1900’s) • 2 electrodes in a glass tube with a gas at low pressure • voltage applied to tube causing a glow discharge • “rays” emitted from cathode (- end) to anode (+ end) • Cathode Rays must be negatively charged!

  7. Discovery and Properties of Electrons • J.J. Thomson - (1897) - changed cathode ray tube experiments by adding two adjustable voltage electrodes into the experiment

  8. Discovery and Properties of Electrons

  9. Discovery and Properties of Electrons • measured charge to mass ratio of electrons • e/m = -1.75881 x 108 coulomb/g of e- • named cathode rays electrons • Thomson is the “discoverer of electrons”

  10. Discovery and Properties of Electrons • Robert A. Millikan - 1st American Nobel Laureate • determined the charge and mass of the electron (1909) • oil drop experiment

  11. Discovery and Properties of Electrons • charge on 1 electron = -1.60219 x 10-19 coulomb • using Thomson’s charge to mass ratio we get that the mass of 1 electron is 9.11 x 10-28 g

  12. Canal Rays and Protons • Goldstein (1886) - “Canal Rays” • streams of positively charged particles in cathode rays • flow in opposite direction of cathode rays • must be positive • postulated existence of “proton”

  13. Radioactivity • Spontaneous emission of high energy radiation • A radioactive substance is placed in a shield containing a small hole so that a beam of radiation is emitted from the hole. • The radiation is passed between two electrically charged plates and detected. • Three spots are noted on the detector: • a spot in the direction of the positive plate, • a spot which is not affected by the electric field, • a spot in the direction of the negative plate.

  14. Radioactivity

  15. Radioactivity • A high deflection towards the positive plate corresponds to radiation which is negatively charged and of low mass. This is called b-radiation (consists of electrons). • No deflection corresponds to neutral radiation. This is called g-radiation. • Small deflection towards the negatively charged plate corresponds to high mass, positively charged radiation. This is called a-radiation.

  16. Discovery of the Nuclear Atom • Plum-Pudding? • Thomson assumed all these charged species were found in a sphere.

  17. Rutherford and the Nuclear Atom • Ernest Rutherford - 1910 - basic picture of atom Geiger & Marsden’s experiment on a- particle scattering from thin Au foils

  18. Rutherford and the Nuclear Atom • In order to get the majority of -particles through a piece of foil to be undeflected, the majority of the atom must consist of a low mass, diffuse negative charge - the electron. • To account for the small number of high deflections of the -particles, the center or nucleus of the atom must consist of a dense positive charge.

  19. Rutherford and the Nuclear Atom • Rutherford decoded the scattering information

  20. Rutherford and the Nuclear Atom • atom is mostly empty space • very small, dense center called nucleus • nearly all of atom’s mass in nucleus • nuclear diameter is 1/10,000 to 1/100,000 times less than atom’s radius • nuclear density is 1015g/mL • equivalent to 3.72 x 109 tons/in3

  21. Neutrons • James Chadwick - 1932 • analyzed evidence from a-particle scattering off Be • recognized existence of massive neutral particles - “neutrons”

  22. The Modern View of Atomic Structure • The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons). • Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is due to the nucleus. • There can be a variable number of neutrons for the same number of protons. Isotopes have the same number of protons but different numbers of neutrons. • Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons.

  23. The Modern View of Atomic Structure

  24. Fundamental Particles • Three fundamental particles make up atoms. The following table lists these particles together with their masses and their charges.

  25. Mass Number & Isotopes • H.G. J. Moseley (1912-1914) - recognized that atomic number is the defining difference between elements • new understanding of Mendeleev’s periodic law

  26. Atomic Number • Sometimes given the symbol Z • number of protons in the nucleus • determines the element • also determines number of electrons in a neutral atom

  27. Isotopes, Atomic Numbers & Mass Numbers • All atoms of an element have the same number of protons in the nucleus • Isotopes of an atom have a different number of neutrons in the nucleus • Atomic number = # of protons • Mass number = # protons + # neutrons

  28. Isotopes, Atomic Numbers & Mass Numbers • By convention, for element X, we write • Isotopes have the same Z but different A.

  29. Isotopes, Atomic Numbers & Mass Numbers

  30. Isotopes • Give the number of protons, neutrons and electrons in each of the following species: 56Fe 56Fe3+31P 31P3-

  31. Atomic Weights • weighted average of the masses of the constituent isotopes • lower number on periodic chart • How do we know what the values of these numbers are?

  32. The Periodic Table • The Periodic Table is used to organize the 114 elements in a meaningful way. • As a consequence of this organization, there are periodic properties associated with the periodic table.

  33. The Periodic Table • Columns in the periodic table are called groups (numbered from 1A to 8A or 1 to 18). • Rows in the periodic table are called periods. .

  34. The Periodic Table • Metals are located on the left hand side of the periodic table (most of the elements are metals). • Non-metals are located in the top right hand side of the periodic table. • Elements with properties similar to both metals and non-metals are called metalloids and are located at the interface between the metals and non-metals.

  35. The Periodic Table

  36. The Periodic Table • Some of the groups in the periodic table are given special names. • These names indicate the similarities between group members: Group 1A: Alkali metals. Group 2A: Alkaline earth metals. Group 6A: Chalcogens. Group 7A: Halogens. Group 8A: Noble gases.

  37. The Periodic Table • Name the following elements. Indicate if each is a metal, nonmetal, or metalloid. • P Sn Mn • K Cu Hg • F As N • Si Na Ca • Fe Ag Mg

  38. Molecules and Molecular Compounds • Molecules are assemblies of two or more atoms bonded together. • Each molecule has a chemical formula. • The chemical formula indicates • which atoms are found in the molecule • in what proportion they are found. • Compounds formed from molecules are molecular compounds.

  39. Molecules and Molecular Compounds • empirical formula - simplest molecular formula, shows ratios of elements but not actual numbers of elements • molecular formula - actual numbers of atoms of each element in the compound

  40. Molecules and Molecular Compounds • Molecules occupy three dimensional space. • However, we often represent them in two dimensions. • The structural formula gives the connectivity between individual atoms in the molecule. • The structural formula may or may not be used to show the three dimensional shape of the molecule.

  41. Molecules and Molecular Compounds

  42. Molecules and Molecular Compounds • If the structural formula does show the shape of the molecule, then either a perspective drawing, ball-and-stick model, or space-filling model is used.

  43. Chemical Formulas • show the ratio of the elements present in the molecule or compound • He, Au, Na - monatomic • O2, H2, Cl2 - diatomic • O3, S4, P8 - more complex elements • H2O, C12H22O11 - compounds

  44. Ions & Ionic Compounds • ions are atoms or groups of atoms that are charged • two basic types of ions • positive ions or cations • one or more electrons less than neutral • negative ions or anions • one or more electrons more than neutral

  45. Ions and Ionic Compounds • The number of electrons an atom loses is related to its position on the periodic table. • Metals tend to form cations whereas non-metals tend to form anions.

  46. Ions and Ionic Compounds • The majority of chemistry involves the transfer of electrons between species. (Ionic Bonding) • Example: • To form NaCl, the neutral sodium atom, Na, must lose an electron to become a cation: Na+. • The electron cannot be lost entirely, so it is transferred to a chlorine atom, Cl, which then becomes an anion: Cl-. • The Na+ and Cl- ions are attracted to form an ionic NaCl lattice which crystallizes.

  47. Ions & Ionic Compounds • Sodium chloride - table salt is an ionic compound

  48. Ion Names and Formulas • Common Polyatomic Ions • Table 2.4 & 2.5

  49. Ion Names and Formulas

  50. Ion Names and Formulas

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