1 / 115

Chapter 2

Chapter 2. Atoms, Molecules, and Ions. Chapter 2: Topics. Early history of chemistry Fundamental chemical laws Dalton’s atomic theory Early experiments to characterize the atom The modern view of atomic structure Molecules and ions An introduction to the Periodic Table

elaine
Download Presentation

Chapter 2

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. Chapter 2 Atoms, Molecules, and Ions

  2. Chapter 2: Topics • Early history of chemistry • Fundamental chemical laws • Dalton’s atomic theory • Early experiments to characterize the atom • The modern view of atomic structure • Molecules and ions • An introduction to the Periodic Table • Naming simple compounds

  3. 2.1 The early history of chemistry • Greeks • Democritus and others - atomos • Alchemy • 1660 - Robert Boyle- experimental definition of element. • Lavoisier- Father of modern chemistry.

  4. Greeks • Matter is composed of fire, earth, water and air • The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”) • He believed that atoms were indivisible and indestructible • His ideas did agree with later scientific theory, but did not explain chemical behavior, and was not based on the scientific method– but just philosophy

  5. Alchemy • Turning Cheep metals into gold • Alchemists discovered several elements and prepared mineral acids

  6. 17th Century • Robert Boyle: First “chemist” to perform quantitative experiments • He published his first book: “The Skeptical Chemist” in 1661. • He talked about elements

  7. 18th Century • George Stahl: Phlogiston flows out of a burning material. • Joseph Priestley: Discovers oxygen gas, “dephlogisticated air, i.e., low in phlogistone”

  8. 2.2 Fundamentals chemical Laws • Law of Conservation of Mass • Law of Definite Proportion • Law of Multiple

  9. Law of Conservation of Mass • It was discovered by Antoine Lavoisier • It was the basis for development of chemistry in the 19th century • Mass is neither created nor destroyed • Combustion involves oxygen, not phlogiston

  10. Law of Definite Proportion(Proust’s Law) • A given compound always contains exactly the same proportion of elements by mass. • Water is composed of 11.1% H and 88.9% O (w/w)

  11. Law of Multiple Proportions • When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers. • The ratio of the masses of oxygen that combine with 1g of H in H2O and H2O2 will be a small whole number (“2”).

  12. Example • Water, H2O has 8 g of oxygen per 1g of hydrogen. • Hydrogen peroxide, H2O2, has 16 g of oxygen per 1g of hydrogen. • 16/8 = 2/1 • Small whole number ratios. • This fact could be explained in terms of atoms

  13. 2.3 Dalton’s Atomic Theory (1808) • Elements are made up of small particles called atoms • Atoms of eachelement are identical. Atoms of different elements are different. • Compounds are formed when atoms combine. Each compound has a always same type and relative number of atoms • Chemical reactions are rearrangement of atoms but atoms are never changed into atoms of other element. , or created or destroyed.

  14. Gay-Lussac hypothesis (1809) • Provided basics to determining absolute formulas of compounds • Gay-Lussac- under the same conditions of temperature and pressure, compounds always react in whole number ratios by volume. • 2volumes of H react with one volume of O to form 2volumes of gaseous water and

  15. Avogadro’s Hypothesis (1811) • 5 liters of oxygen • 5 liters of nitrogen • Same number of particles! At the same temperature and pressure, equal volumes of different gases contain the same number of particles.

  16. If Avogadro's hypothesis is correct, Gay-Lussac’s can be interpreted as follows: • 2 molecules of H react with 1 molecule of O 2 molecules of H2O

  17. 2.4 Early experiments to characterize the atom • Based on Dalton, Gay-Lussac, Avogadro and others, work started to identify the nature of the atom • What is an atom made of? How do atoms of various elements differ?

  18. The electron • J. J. Thomson - postulated the existence of electrons using cathode ray tubes. • Ernest Rutherford - explained the nuclear atom, containing a dense nucleus with electrons traveling around the nucleus at a large distance.

  19. Thomson’s Experiment Voltage source - + When high voltage is applied to the tube a ray emanates from the cathode is called cathode ray.

  20. Voltage source Thomson’s Experiment - +

  21. Voltage source Thomson’s Experiment - + • Passing an electric current makes a beam appear to move from the negative to the positive end.

  22. Voltage source Thomson’s Experiment • By adding an electric field

  23. Voltage source Thomson’s Experiment + - • By adding an electric field, he found that the moving particles were negatively charged

  24. Results of Thomson Experiment • Electrons are produced from electrodes made from various types of metals, all atoms must contain electrons. • Since atoms are electrically neutral, they must contain positively chargedparticles. • Thomson determined charge-to-mass ratio of an electron: • e/m = -1.76X108C/g

  25. Thomson’s Model • Atom consisted of a diffuse cloud of positive charge with negative electrons embedded randomly • Atom was like plum pudding. • Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

  26. Atomizer Oil droplets + - Oil Telescope Millikan’s Experiment

  27. Millikan’s Experiment X-rays X-rays give some electrons a charge.

  28. Millikan’s Experiment From the mass of the drop and the charge on the plates, the mass of an electron is calculated

  29. Radioactivity • Certain elements produce high energy radiation • Discovered by accident and was a result of spontaneous emission by uranium • Bequerel (1896) found that a piece of mineral containing uranium could produce an image on a photographic plate in the absence of light. • Three types of radiation were known: • alpha- helium nucleus (+2 charge, 7300 times that of the electron) • beta- high speed electron • gamma- high energy light

  30. The nuclear atomRutherford’s Experiment • Aimed at testing Thomson’s plum pudding model • Used uranium to produce alpha particles. • Alpha particles are directed at gold foil through hole in lead block. • Since the mass is evenly distributed in gold atoms alpha particles should go straight through. • Used gold foil because it could be made atoms thin.

  31. Florescent Screen Lead block Uranium Gold Foil

  32. What he expected

  33. Because

  34. Because, he thought the mass was evenly distributed in the atom.

  35. What he got

  36. + How he explained it • Atom is mostly empty • Small dense, positive particle at center. • Alpha particlesare deflected by it if they get close enough.

  37. + Proof for nuclear atom

  38. Nuclear atom model • According to Rutherford: The atom consists of a dense center of positive charge (Nucleus) with electrons moving around it at distance that is large relative to the nuclear radius

  39. 2.5 The modern view of an atomic structure:An introduction • The atom is mostly empty space. • Two regions • Nucleus- protons and neutrons. • It is characterized by small size and high density • Electron cloud- region where you might find an electron. • The chemistry of atom • Results mainly from electrons

  40. A cross section of nuclear atom

  41. Mass and charge of nuclear particles Particle Mass (Kg) Charge Electron 9.11X10-31 -1 Proton 1.67X10-27 +1 Neutron 1.67X10-27 None

  42. Why atoms of different elements have different properties? • Atoms of different elements have different number of protons and electrons. • Number and arrangement of electrons around nucleus differ from one element to another.

  43. Sub-atomic Particles • Z - atomic number = number of protons determines type of atom. • A - mass number = number of protons + neutrons. • Number of protons = number of electrons if atom is neutral.

  44. Symbols A Mass number X Atomic number Z 23 Na 11 Na-23

  45. More Atomic Symbols 16 31 65 O P Zn 8 15 30 8 p+ 15 p+ 30 p+ 8 n 16 n 35 n 8 e- 15e- 30 e-

  46. Isotopes • Atoms of the same element (same atomic number) with different mass numbers • Atoms with the same number of protons, but different numbers of neutrons. Isotopes of chlorine 35Cl 37Cl 1717 chlorine - 35 chlorine – 37 Cl-35 Cl-37

  47. Two isotopes of sodium • Isotopes show almost identical chemical properties. Why? • They possess same number of electrons

  48. 2.6 Molecules and ionsIntroduction to chemical bonding • The forces that hold atoms together are called chemical bonding • Covalent bonding - sharing electrons. • Collection of atoms by covalent bonding lead to molecules • Molecules can be represented by formulas • Chemical formula- Symbol relates number and type of atoms in a molecule. • Diatomic molecule: two atoms of same element are connected by a covalent bond.

  49. Molecules that contain two atoms of the same element bonded together are called diatomic molecules.

  50. Molecular Compounds • Molecular formulas • give the actual numbers and types of atoms in a molecule. • Examples: H2O, CO2, CO, CH4, H2O2, O2, O3, and C2H4. • Structural formula: bonds are shown as lines H H H C C H H H

More Related