1 / 49

Atoms, Ions and Molecules The Building Blocks of Matter

Atoms, Ions and Molecules The Building Blocks of Matter. Chapter 2. Chapter Outline. 2.1 The Rutherford Model of Atomic Structure 2.2 Nuclides and Their Symbols 2.3 Navigating the Periodic Table 2.4 The Masses of Atoms, Ions, and Molecules 2.5 Moles and Molar Mass 2.6 Making Elements

avel
Download Presentation

Atoms, Ions and Molecules The Building Blocks of Matter

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Atoms, Ions and MoleculesThe Building Blocks of Matter Chapter 2

  2. Chapter Outline 2.1 The Rutherford Model of Atomic Structure 2.2Nuclides and Their Symbols 2.3 Navigating the Periodic Table 2.4 The Masses of Atoms, Ions, and Molecules 2.5 Moles and Molar Mass 2.6 Making Elements 2.7 Artificial Nuclides

  3. Experiments in Atomic Structure • J. J. Thompson (1906 Nobel Prize in Physics)- cathode ray tube experiments; discovery of the electron; measurement of the charge-to-mass ratio. • Robert Millikan (1923 Nobel Prize in Physics) - oil-drop experiments; measured the mass of the electron, therefore calculate the charge • Ernest Rutherford (1908 Nobel Prize in Physics) - gold-foil experiments; the nuclear atom • James Chadwick (1935 Nobel Prize in Physics) - discovery of the neutron

  4. J.J. Thomson Cathode Ray Tube Experiments - Electrons

  5. Results of “Cathode Ray” Experiments • Travel in straight lines • invisible • independent of cathode composition • bend in a magnetic field like a negatively-charged particle would • charge/mass = -1.76 x 108 C/g

  6. Thompson’s “Plum Pudding” Model of the Atom electrons distributed throughout a diffuse, positively charged sphere.

  7. Robert Millikan’s oil drop Experiment - measured the mass of the electron

  8. Millikan’s Results • The air molecules in the chamber were ionized by a beam of X-rays, producing electrons and positively-charged fragments • Fine mist of oil introduced into chamber; electrons adhere to the droplets • Negatively-charged droplets settle to bottom of chamber under influence of gravity • Charged repeller plates adjusted until droplets were suspended in mid-air • From the physics and knowledge of the size of the gravitational and electrostatic forces, the charge on each droplet could be calculated • Discovered that each droplet was a whole-number multiple of 1.60 X 10-19 C, so the mass = 9.11 X 10-28 g

  9. Radioactivity and the Nuclear Atom Spontaneous emission of particles and/or radiation from a decaying, unstable nucleus -particles = -particles = -rays =

  10. Ernest Rutherford - the nuclear atom

  11. Rutherford's Observations b) Expected results from “plum pudding” model. c) Actual results. • the majority of particles penetrated undeflected • some particles were deflected at small angles • occasionally -particles scattered back at large angles

  12. Rutherford’s Conclusions • The atom is mainly empty space because most of the -particles passed through undeflected • The nucleus is very dense and positively charged because some of the -particles were repulsed and deflected • Electrons occupy the space around the nucleus • The atom is electrically neutral

  13. Rutherford’s Model of the Atom atomic radius ~ 100 pm = 1 x 10-10 m nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m If the nucleus was the size of an orange, then the radius of the atom would be 2.5 miles

  14. mass p  mass n = 1840 x mass e-

  15. Chapter Outline 2.1 The Rutherford Model of Atomic Structure 2.2Nuclides and Their Symbols 2.3 Navigating the Periodic Table 2.4 The Masses of Atoms, Ions, and Molecules 2.5 Moles and Molar Mass 2.6 Making Elements 2.7 Artificial Nuclides

  16. Atomic Mass Units • Atomic Mass Units (amu) • Comprise a relative scale to express the masses of atoms and subatomic particles. • Scale is based on the mass of 1 atom of carbon: • 6 protons + 6 neutrons = 12 amu. • 1 amu = 1 Dalton (Da)

  17. Isotopes: Experimental Evidence

  18. A X Mass Number Element Symbol Z Atomic Number 1 3 2 H (T) H (D) H 1 1 1 238 235 U U 92 92 Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes (nuclides)are atoms of the same element with different numbers of neutrons in the nucleus

  19. Practice: Isotopic Symbols • Use the format AX to write the symbol for the nuclides having 28 protons and 31 neutrons. • Collect and Organize: • Analyze: • Solve: • Think about It:

  20. Practice: Identifying Atoms and Ions • Complete the missing information in the table. • Collect and Organize: • Analyze: • Solve: • Think about It:

  21. Chapter Outline 2.1 The Rutherford Model of Atomic Structure 2.2Nuclides and Their Symbols 2.3 Navigating the Periodic Table 2.4 The Masses of Atoms, Ions, and Molecules 2.5 Moles and Molar Mass 2.6 Making Elements 2.7 Artificial Nuclides

  22. The Periodic Table of the Elements Mendeleev’s Periodic Table Dmitrii Mendeleev (1872): • Ordered elements by atomic mass. • Arranged elements in columns based on similar chemical and physical properties. • Left open spaces in the table for elements not yet discovered.

  23. The Modern Periodic Table • Also based on a classification of elements in terms of their physical and chemical properties. • Horizontal rows: called periods (1 → 7). • Columns: contain elements of the same family or group (1 →18). • Several groups have names as well as numbers.

  24. Navigating the Modern Periodic Table – Groups and Families

  25. Groups of Elements (cont.)

  26. These 7 elements occur naturally as diatomics (memorize) - H2 N2 F2 O2 I2 Cl2 Br2

  27. Metals • found to the left of the “diagonal line” • lose electrons in chemical reactions • solids (except for Hg, Cs, and Fr) • conduct electricity • ductile (draw into a wire) • malleable (roll into sheets) • form alloys ("solid-solution" of one metal in another)

  28. Nonmetals • found to the right of the “diagonal line” • like to gain electrons from metals, or share electrons among themselves • found as solids, liquids (Br), and gases (Inert gases, and H, N, O, F, Cl) • “diatomics” - H2, N2, F2, O2 ,I2, Cl2, Br2 • oxygen also exist as ozone, O3 • insulators (except for graphite or C) Helium-Neon lasers

  29. Metalloids • elements next to the “diagonal line” • B, Si, Ge, As, Sb, and Te • physical properties of a metal (can be “convinced” to conduct electricity) and chemical properties of a nonmetal Elemental Si is used in the semiconductor industry

  30. Chapter Outline 2.1 The Rutherford Model of Atomic Structure 2.2Nuclides and Their Symbols 2.3 Navigating the Periodic Table 2.4 The Masses of Atoms, Ions, and Molecules 2.5 Moles and Molar Mass 2.6 Making Elements 2.7 Artificial Nuclides

  31. Average Atomic Mass • Weighted average mass of natural sample of an element, calculated by multiplying the natural abundance of each isotope by its exact mass in amu’s and then summing up these products. AM = (mass 1)(abn) + (mass 2)(abn) + (mass 3)(abn) +………

  32. Molecular Mass Molecular Mass – the sum of the average atomic masses of the atoms in it. e.g. H2SO4 NOTE: the terms mass and weight are used interchangeably, e.g. molecular weight (MW) or atomic weight (AW)

  33. Formula Units and Formula Mass Formula Units – for ionic compounds, the smallest electrically neutral unit in an ionic compound Formula Mass – the sum of the average atomic masses of the cations and anions that make up a neutral formula unit e.g. NaCl

  34. Chapter Outline 2.1 The Rutherford Model of Atomic Structure 2.2Nuclides and Their Symbols 2.3 Navigating the Periodic Table 2.4 The Masses of Atoms, Ions, and Molecules 2.5 Moles and Molar Mass 2.6 Making Elements 2.7 Artificial Nuclides

  35. The Mole - The mole is the Chemist’s counting unit dozen = 12 gross = 144 pair = 2 Avogadro’s Number (NA) = 6.022 X 1023 = 1 mole of atoms, molecules, ions, etc. ream = 500

  36. One Mole of: S C Hg Fe Cu

  37. Experiment – how many atoms must be added together so that the mass in grams = mass in amu’s? Analogy using coins: Mass ratio = 1 : 5 : 25

  38. Mass in amu’s Mass in grams/mole Significance of the Mole Equivalent to NA of carbon atoms weighs __________ NA of iron atoms weighs __________

  39. Moles, Mass, and Particles • To convert between number of particles and an equivalent number of moles.

  40. Sample Exercise 2.5 The silicon used to make computer chips has to be extremely pure. Fpr example, it must contain less than 3 x 10-10 moles of phosphorus (a common impurity in Si) per mole of silicon. What is this level of impurity expressed in atoms of phosphorus per mole of Si?

  41. Using the Molar Mass as a Conversion Factor for Atoms & Molecules e.g. carbon e.g. H2SO4 sulfuric acid

  42. Moles, Mass, and Particles grams of atoms or molecules moles of atoms or molecules Numbers of atoms or molecules

  43. Practice: Mole Calculations #1 • How many moles of K atoms are present in 19.5 g of potassium? • How many atoms of K are there?

  44. Practice: Mole Calculations #2 How many moles are present in 58.4 g of chalk (CaCO3)?

  45. Practice: Mole Calculations #3 The uranium used in nuclear fuel exists in nature in several minerals. Calculate how many moles of uranium are found in 100.0 grams of carnotite, K2(UO2)2(VO4)2•3H2O.

  46. Practice: Mole Calculations #4 Convert 2.45 x 1018 molecules of KCl to grams

More Related