Chapter 5 – Early Atomic Theory & Structure

1. Chapter 5 – Early Atomic Theory & Structure 5.1 Early Thoughts 5.6 Subatomic Parts of the Atom 5.7 The Nuclear Atom 5.2 Dalton's Model of the Atom 5.8 General Arrangement ofSubatomic Particles 5.3 Composition of Compounds 5.9 Atomic Numbers of theElements 5.4 The Nature of Electric Charge 5.5 Discovery of Ions 5.10 Isotopes of the Elements 5.11 Atomic Mass

2. Early Thoughts • The earliest models of the atom were developed by the ancient Greek philosophers. • Empedocles stated that matter was made of 4 elements: earth, air, fire, and water. • Democritus (about 470-370 B.C.) thought that all forms of matter were divisible into tiny indivisible particles. He called them “atoms” from the Greek “atomos” indivisible.

3. Early Thoughts • Aristotle (384-322 B.C.) rejected the theory of Democritus and advanced the Empedoclean theory. • Aristotle’s influence dominated the thinking of scientists and philosophers until the beginning of the 17th century.

4. Dalton’s Model of the Atom 2000 years after Aristotle, John Dalton an English schoolmaster, proposed his model of the atom–which was based on experimentation.

5. Modern research has demonstrated that atoms are composed of subatomic particles. Atoms under special circumstances can be decomposed. Dalton’s Atomic Theory • Elements are composed of minute indivisible particles called atoms. • Atoms of the same element are alike in mass and size. • Atoms of different elements have different masses and sizes. • Chemical compounds are formed by the union of two or atoms of different elements.

6. Dalton’s Atomic Theory • Atoms combine to form compounds in simple numerical ratios, such as one to one , two to two, two to three, and so on. • Atoms of two elements may combine in different ratios to form more than one compound.

7. Dalton’s atoms were individual particles. Atoms of each element are alike in mass and size. 5.1

8. Dalton’s atoms were individual particles. Atoms of different elements are not alike in mass and size. 5.1

9. Daltons atoms combine in specific ratios to form compounds.

10. Composition of Compounds The Law of Definite Composition A compound always contains two or more elements combined in a definite proportion by mass.

11. Composition of Water • Water always contains the same two elements: hydrogen and oxygen. • The percent by mass of hydrogen in water is 11.2%. • The percent by mass of oxygen in water is 88.8%. • Water always has these percentages. If the percentages were different the compound would not be water.

12. Composition of Compounds The Law of Multiple Proportions Atoms of two or more elements may combine in different ratios to produce more than one compound.

13. Composition of Hydrogen Peroxide • Hydrogen peroxide always contains the same two elements: hydrogen and oxygen. • The percent by mass of hydrogen in hydrogen peroxide is 5.9%. • The percent by mass of oxygen in hydrogen peroxide is 94.1%. • Hydrogen peroxide always has these percentages. If the percentages were different the compound would not be hydrogen peroxide.

14. Combining Ratios of Hydrogen and Oxygen • The formula for water is H2O. • The formula for hydrogen peroxide is H2O2. • Hydrogen peroxide has twice as many oxygens per hydrogen atom as does water.

15. Composition of Compounds

16. q1 and q2 are charges, r is the distance between charges and k is a constant. The Nature of Electric Charge Properties of Electric Charge • Charge may be of two types: positive and negative. • Unlike charges attract (positive attracts negative), and like charges repel (negative repels negative and positive repels positive). • Charge may be transferred from one object to another, by contact or induction. • The less the distance between two charges, the greater the force of attraction between unlike charges (or repulsion between identical charges).

17. Discovery of Ions • Michael Faraday discovered that certain substances when dissolved in water conducted an electric current. • He found that atoms of some elements moved to the cathode (negative electrode) and some moved to the anode (positive electrode). • He concluded they were electrically charged and called them ions (Greek wanderer).

18. Δ NaCl → Na+ + Cl- Discovery of Ions • Arrhenius accounted for the electrical conduction of molten sodium chloride (NaCl) by proposing that melted NaCl dissociated into the charged ions Na+ and Cl-. • Svante Arrhenius reasoned that an ion is an atom (or a group of atoms) carrying a positive or negative electric charge.

19. Discovery of Ions • In the melt the positive Na+ ions moved to the cathode (negative electrode). Thus positive ions are called cations. • In the melt the negative Cl- ions moved to the anode (positive electrode). Thus negative ions are called anions. NaCl → Na+ + Cl-

20. Subatomic Parts of the Atom The diameter of an atom is 0.1 to 0.5 nm. This is 1 to 5 ten billionths of a meter. If the diameter of this dot is 1 mm, then 10 million hydrogen atoms would form a line across the dot. Even smaller particles than atoms exist. These are called subatomic particles.

21. Subatomic Particle - Electron • Crookes tubes experiments led the way to an understanding of the subatomic structure of the atom. • Crookes tube emissions are called cathode rays. • In 1875 Sir William Crookes invented the Crookes tube.

22. Subatomic Particle - Electron In 1897 Sir Joseph Thompson demonstrated that cathode rays: • travel in straight lines. • are negative in charge. • are deflected by electric and magnetic fields. • produce sharp shadows • are capable of moving a small paddle wheel.

23. Subatomic Particle - Electron This was the discovery of the fundamental unit of charge – the electron.

24. Subatomic Particle - Proton • Eugen Goldstein, a German physicist, first observed protons in 1886: • Thompson determined the proton’s characteristics. • Thompson showed that atoms contained both positive and negative charges. • This disproved the Dalton model of the atom which held that atoms were indivisible.

25. Subatomic Particle - Neutron • Its actual mass is slightly greater than the mass of a proton. • James Chadwick discovered the neutron in 1932.

26. Subatomic Particles

27. Ions • Negative ions were explained by assuming that extra electrons can be added to atoms. • Positive ions were explained by assuming that a neutral atom loses electrons.

28. When one or more electrons are lost from an atom, a cation is formed. 5.4

29. When one or more electrons are added to a neutral atom, an anion is formed. 5.4

30. The Nuclear Atom • Radioactive elements spontaneously emit alpha particles, beta particles and gamma rays from their nuclei. • By 1907 Rutherford found that alpha particles emitted by certain radioactive elements were helium nuclei. • Radioactivity was discovered by Becquerel in 1896.

31. The Rutherford Experiment • Most of the alpha particles passed through the foil with little or no deflection. • He found that a few were deflected at large angles and some alpha particles even bounced back. • Rutherford in 1911 performed experiments that shot a stream of alpha particles at a gold foil.

32. The Rutherford Experiment Rutherford’s alpha particle scattering experiment. 5.5

33. The Rutherford Experiment • An electron with a mass of 1/1837 amu could not have deflected an alpha particle with a mass of 4 amu. • Rutherford knew that like charges repel. • Rutherford concluded that each gold atom contained a positively charged mass that occupied a tiny volume. He called this mass the nucleus.

34. The Rutherford Experiment • Most of the alpha particles passed through the gold foil. This led Rutherford to conclude that a gold atom was mostly empty space. • If a positive alpha particle approached close enough to the positive mass it was deflected.

35. The Rutherford Experiment • Because alpha particles have relatively high masses, the extent of the deflections led Rutherford to conclude that the nucleus was very heavy and dense.

36. The Rutherford Experiment Deflection Scattering Deflection and scattering of alpha particles by positive gold nuclei. 5.5

37. General Arrangement of Subatomic Particles • Rutherford’s experiment showed that an atom had a dense, positively charged nucleus. • Chadwick’s work in 1932 demonstrated the atom contains neutrons. • Rutherford also noted that light, negatively charged electrons were present in an atom and offset the positive nuclear charge.

38. General Arrangement of Subatomic Particles • The negative electrons surround the nucleus. • The nucleus contains protons and neutrons • Rutherford put forward a model of the atom in which a dense, positively charged nucleus is located at the atom’s center.

39. General Arrangement of Subatomic Particles 5.6

40. Atomic Numbers of the Elements • The atomic number of an atom determines which element the atom is. • The atomic number of an element is equal to the number ofprotons in the nucleus of that element.

41. Atomic Numbers of the Elements Every atom with an atomic number of 1 is a hydrogen atom.Every hydrogen atom contains 1 proton in its nucleus.

42. atomic number 1 proton in the nucleus 1H Every atom with an atomic number of 1 is a hydrogen atom.

43. Atomic Numbers of the Elements Every atom with an atomic number of 6 is a carbonatom.Every carbon atom contains 6 protons in its nucleus.

44. atomic number 6 protons in the nucleus 6C Every atom with an atomic number of 6 is a carbon atom.

45. atomic number 92 protons in the nucleus 92U Every atom with an atomic number of 92 is a uranium atom.

46. Atomic Numbers of the Elements • They always have the same number of protons, but they can have different numbers of neutrons in their nuclei. • The difference in the number of neutrons accounts for the difference in mass. • These are isotopes of the same element. • Atoms of the same element can have different masses.

47. Atomic Numbers of the Elements Isotopes of the Same Element Have Equal numbers of protons Different numbers of neutrons

48. Isotopic Notation

49. 12 C 6 Isotopic Notation 6 protons + 6 neutrons 6 protons

50. Isotopic Notation 6 protons + 8 neutrons 14 C 6 6 protons