Chapter 3

1. Atoms: The Building Blocks of Matter Chapter 3

2. Section 1 The particle theory of matter was supported as early as 400 b.c. by certain Greek thinkers, such as Democritus. He called nature’s basic particle an atom, based on the Greek word meaning “indivisible.” Aristotle was part of the generation that succeeded Democritus. His ideas had a lasting impact on Western civilization, and he did not believe in atoms. He thought that all matter was continuous, and his opinion was accepted for nearly 2000 years. Neither the view of Aristotle nor that of Democritus was supported by experimental evidence, so each remained speculation until the eighteenth century. Then scientists began to gather evidence favoring the atomic theory of matter.

3. Foundations of Atomic Theory • Nearly all chemists in late 1700s accepted the definition of an element as a substance that cannot be broken down further • Knew about chemical reactions • Great disagreement as to whether elements always combine in the same ratio when forming a specific compound

4. Law of Conservation of Mass • With the help of improved balances, investigators could accurately measure the masses of the elements and compounds they were studying • This lead to discovery of several basic laws • Law of conservation of mass states that mass is neither destroyed nor created during ordinary chemical reactions or physical changes

6. Law of Definite Proportions • Was quickly followed by the 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 Each of the salt crystals shown here contains exactly 39.34% sodium and 60.66% chlorine by mass.

8. Law of Multiple Proportions • Two elements sometimes combine to form more than one compound • For example, the elements carbon and oxygen form two compounds, carbon dioxide and carbon monoxide • Consider samples of each of these compounds, each containing 1.0 g of carbon • In carbon dioxide, 2.66 g of oxygen combine with 1.0 g of carbon • In carbon monoxide, 1.33 g of oxygen combine with 1.0 g of carbon • The ratio of the masses of oxygen in these two compounds is exactly 2.66 to 1.33, or 2 to 1

10. 1808 John Dalton • Proposed an explanation for the law of conservation of mass, the law of definite proportions, and the law of multiple proportions • He reasoned that elements were composed of atoms and that only whole numbers of atoms can combine to form compounds • His theory can be summed up by the following statements

11. Dalton’s Atomic Theory • All matter is composed of extremely small particles called atoms. • Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. • Atoms cannot be divided, created or destroyed. • Atoms of different elements combine in simple whole-number ratios to form chemical compounds. • In chemical reactions, atoms are combined, separated, or rearranged.

13. Conservation of Mass Explained • Chemical reactions involve the combination, separation, or rearrangement of atoms • During these processes atoms are not divided, created, or destroyed

14. Definite Proportions Explained • A given chemical compound is always made of the same combination of atoms

15. Multiple Proportions Explained • Example of carbon oxides, the ratio of 2-to-1 ratio of oxygen masses results because carbon dioxide always contains twice as many atoms of oxygen (per atom of carbon) as does carbon monoxide

16. Modern Atomic Theory • Dalton turned Democritus’sidea into a scientific theory which was testable • Not all parts of his theory have been proven correct • Ex. We know atoms are divisible into even smaller particles • We know an element can have atoms with different masses

17. Section 2 The Structure of the Atom

18. All atoms consist of two regions • Nucleus very small region located near the center of an atom • In the nucleus there is at least one positively charged particle called the proton • Usually at least one neutral particle called the neutron • Surrounding the nucleus is a region occupied by negatively charged particles called electrons

20. Discovery of the Electron • Resulted from investigations into the relationship between electricity and matter • Late 1800s, many experiments were performed  electric current was passed through different gases at low pressures • These experiments were carried out in glass tubes known as cathode-ray tubes

21. Cathode Rays and Electrons • Investigators noticed that when current was passed through a cathode-ray tube, the opposite end of the tube glowed • Hypothesized that the glow was caused by a stream of particles, which they called a cathode ray • The ray traveled from the cathode to the anode when current was passed through the tube

23. Observations • 1. An object placed between the cathode and the opposite end of the tube cast a shadow on the glass • 2. A paddle wheel placed on rails between the electrodes rolled along the rails from the cathode toward the anode • These facts supported the existence of a cathode ray • The paddlewheel experiment showed that a cathode ray had sufficient mass to set the wheel in motion

24. 3. Cathode rays were deflected by a magnetic field in the same way as a wire carrying electric current, which was known to have a negative charge • 4. The rays were deflected away from a negatively charged object

25. A paddle wheel placed in the path of the cathode ray moves away from the cathode and toward the anode. The movement of the wheel led scientists to conclude that cathode rays have mass.

26. Observations led to the hypothesis that the particles that compose cathode rays are negatively charged • Strongly supported by a series of experiments carried out in 1897 by the English physicist Joseph John Thomson

27. He was able to measure the ratio of the charge of cathode-ray particles to their mass • He found that this ratio was always the same, regardless of the metal used to make the cathode or the nature of the gas inside the cathode-ray tube • Thomson concluded that all cathode rays are composed of identical negatively charged particles, which were later named electrons

28. Charge and Mass of the Electron • Thomson’s experiment revealed that the electron has a very large charge for its tiny mass • Mass of the electron is about one two-thousandth the mass of the simplest type of hydrogen atom (the smallest atom known) • Since then found that the electron has a mass of 9.109 × 10−31 kg, or 1/1837 the mass of the hydrogen atom

29. Thomson’s Atom

30. 1909 Robert A. Millikan • Confirmed that the electron carries a negative electric charge • Because cathode rays have identical properties regardless of the element used to produce them, it was concluded that electrons are present in atoms of all elements • Cathode-ray experiments provided evidence that atoms are divisible and that one of the atom’s basic constituents is the negatively charged electron

32. Two more inferences about the atom • 1. Because atoms are electrically neutral, they must contain a positive charge to balance the negative electrons • 2. Because electrons have so much less mass than atoms, atoms must contain other particles that account for most of their mass

33. Discovery of Atomic Nucleus • 1911 by New Zealander Ernest Rutherford and his associates Hans Geiger and Ernest Marsden • Bombarded a thin, gold foil with fast-moving alpha particles (positively charged particles with about four times the mass of a hydrogen atom) • Assume mass and charge were uniformly distributed throughout atoms of gold foil (from Thomson’s model of the atom) • Expected alpha particles to pass through with only slight deflection

34. What Really Happened… • Most particles passed with only slight deflection • However, 1/8,000 were found to have a wide deflection • Rutherford explained later it was “as if you have fired a 15-inch artillery shell at a piece of tissue paper and it came back and hit you.”

37. Explanation • After 2 years, Rutherford finally came up with an explanation • The rebounded alpha particles must have experienced some powerful force within the atom • The source of this force must occupy a very small amount of space because so few of the total number of alpha particles had been affected by it • The force must be caused by a very densely packed bundle of matter with a positive electric charge • Rutherford called this positive bundle of matter the nucleus

38. Rutherford had discovered that the volume of a nucleus was very small compared with the total volume of an atom • If the nucleus were the size of a marble, then the size of the atom would be about the size of a football field • But where were the electrons? • Rutherford suggested that the electrons surrounded the positively charged nucleus like planets around the sun • He could not explain, however, what kept the electrons in motion around the nucleus

39. Rutherford’s Atom

40. Composition of Atomic Nucleus • Except hydrogen, all atomic nuclei made of two kinds of particles • Protons • Neutrons • Protons = positive • Neutrons = neutral • Electrons = negative • Atoms are electrically neutral, so number of protons and electrons IS ALWAYS THE SAME

41. The nuclei of atoms of different elements differ in the number of protons they contain and therefore in the amount of positive charge they possess • Thus, the number of protons in an atom’s nucleus determines that atom’s identity

42. Forces in the Nucleus • Usually, particles that have the same electric charge repel one another • Would expect a nucleus with more than one proton to be unstable • When two protons are extremely close to each other, there is a strong attraction between them • Nuclear forces short-range proton-neutron, proton-proton, and neutron-neutron forces that hold the nuclear particles together

43. Section 3 – Counting Atoms

44. Atomic Number • Atoms of different elements have different numbers of protons • Atomic number (Z) number of protons in the nucleus of each atom of that element • Shown on periodic table • Atomic number identifies an element 3 Li Lithium 6.941 [He]2s1

46. Isotopes • Hydrogen and other atoms can contain different numbers of neutrons • Isotope atoms of same element that have different masses

48. Isotopes of Hydrogen

49. Mass Number • Mass number total number of protons and neutrons in the nucleus of an isotope