Chapter 3

1. Chapter 3 Atoms: The Building Blocks of Matter

2. 3-1 An Ancient Idea • Water: atoms smooth and round so it flowed with no permanent shape • Fire: atoms thorny,making burns painful • Earth: atoms rough and jagged so they held together to make hard stable substances

3. 3-1 Opposing View • Did not believe in atoms • Believed matter was continuous • Very influential • Neither Democritus nor Aristotle supported their ideas with experimentation.

4. 3-1 Law of Conservation of Mass • Antoine Laurent Lavoisier (1743-1794) – the father of modern chemistry • Mass is neither created nor destroyed during ordinary chemical reactions or physical changes • Arrested by French Revolutionary Tribunal for his membership in the Ferme Generale and executed

5. 3-1 Law of Definite Proportions (Constant Composition) • A chemical compound contains the same proportions by mass regardless of the size of the sample or source of the compound • Sodium chloride (NaCl) is always 39.34% sodium and 60.66% chlorine

6. 3-1 John Dalton • English schoolteacher – 1808 • Proposed explanation for these laws

7. 3-1 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 subdivided, 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.

8. 3-1Dalton and Conservation of Mass • In chemical reactions, atoms are rearranged – bonds are broken and new bonds are formed. • Same number of atoms of each type exist before and after the reaction. • No atoms are created or destroyed.

9. 3-1 Dalton and Definite Proportions • Each kind of atom has its own mass and atoms in a compound always exist in fixed whole number ratios. • Sodium chloride always 1:1 sodium:chlorine • Na: 23, Cl: 35.5 • Na: 23/58.5 *100 = 39.34% • Cl: 35.5/58.5 *100 = 60.66%

10. 3-1 Law of Multiple Proportions • If two or more different compounds are composed of the same elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a whole number ratio.

11. 3-1 Dalton and Multiple Proportions • Each kind of atom has a unique mass and atoms in compounds are combined in fixed whole number ratios.

12. 3-1 Impact of Dalton’s Atomic Theory • Related Democritus’s idea of atoms to the MEASURABLE property of mass – atomic theory could then be TESTED by EXPERIMENT • Dalton’s theory has been modified over the years as new information has become available, but the fundamental principles hold true today

13. 3-1 Dalton’s Atomic Model

14. 3-2 Structure of the Atom ATOM - The smallest particle of an element that retains the chemical properties of that element.

15. 3-2 Discovery of the Electron • 1897 • J. J. Thomson, English physicist • Did experiments with cathode ray tubes – glass tubes containing gases at low pressure

16. 3-2 Cathode Ray Tubes • Cathode – negative electrode, anode – positive electrode • When a current is passed through the tube, the end opposite the cathode glows • Glow caused by stream of particles called cathode ray because is originated at cathode – ray traveled from cathode to anode

17. 3-2 CRT Experiments • Observation – a paddle wheel placed on rails between electrodes rolled along rails from cathode to anode • Conclusion – existence of cathode ray supported, cathode ray has MASS

18. 3-2 CRT Experiments • Observations – cathode rays deflected by a magnetic field in the same way as a wire carrying an electric current (cathode ray acts NEGATIVE); deflected away from negatively charged objects • Conclusion – cathode ray is made of negatively charged particles

19. 3-2 Charge and Mass of the Electron • J. J. Thomson called these tiny negatively charged particles “electrons” • He calculated the mass to charge ratio of the electron – it is always the same no matter what metal is used for the electrodes

20. 3-2 Charge and Mass of the Electron • Robert Millikan, American physicist • 1909 • Showed that mass of electron is 1/2000 the mass of the simplest known atom (hydrogen) • 9.109 x 10-31 kg

21. 3-2 Important Points • Atoms contain tiny negatively charged particles called electrons. • Electrons are present in atoms of all elements. • Atoms are divisible and one of parts is negatively charged. • Because atoms are neutral, there must also be a positive component. • Because electrons have such small mass, atoms must contain other parts that make up most of their mass.

22. 3-2 The Plum Pudding Model of the Atom

23. 3-2 Discovery of the Atomic Nucleus • 1911 • Ernest Rutherford, New Zealand (with Hans Geiger and Ernest Marsden) • Important experiment providing more detail into atom’s structure.

24. 3-2 The Gold Foil Experiment • Thin, gold foil bombarded with alpha particles (positively charged particles with 4x mass of hydrogen atom)

25. 3-2 The Gold Foil Experiment - Results • Most particles went right through gold foil • A few were slightly deflected • A few bounced off the gold foil!

26. 3-2 The Gold Foil Experiment - Conclusions • Particles that pass through foil – hit nothing • Particles slightly deflected – come close to a positive charge • Particles that bounce back – hitting a positive charge

27. 3-2 Rutherford’s Explanation • The atom has a densely packed bundle of matter with a positive charge (he called it the nucleus). • The nucleus contains all of the positive charge and most of the mass. • The nucleus has very little volume – the atom is mostly empty space.

28. 3-2 Rutherford’s Atomic Model • Small nucleus in the center • Electrons orbit nucleus like planets around the Sun. • Sometimes called planetary atomic model.

29. 3-2 Composition of the Atomic Nucleus • Nuclei contain two kinds of particles – protons and neutrons • Electrons are outside the nucleus, in the electron cloud • Table 3-1 on p. 74

30. 3-2 Forces in the Nucleus • Like charges generally repel each other BUT at very close range there is an attraction between them • Nuclear Forces – short range proton-neutron, proton-proton, and neutron-neutron forces that hold the nuclear particle together

31. 3-2 The Sizes of Atoms • Expressed as atomic radius – distance from center of nucleus to outer edge of electron cloud • Measured in picometers (10-12 m) • Generally range from 40 pm to 270 pm • Nuclear radius is ~0.001 pm (like a dime in a football field)

32. 3-3 Counting Atoms • Atoms of the same element all have the same number of protons – this is the atomic number • Atomic number is found on the periodic table – elements arranged in order of increasing atomic number

33. 3-3 Isotopes • Atoms of the same element that have different masses (different numbers of neutrons) – same chemical behavior

34. 3-3 Isotopes • Most of the elements consist of mixtures of isotopes. • To identify an isotope, must know atomic number AND mass number (protons + neutrons) • Isotopes usually identified by specifying mass number

35. 3-3 Isotopes Nuclide – general term for any isotope of any element Hyphen Notation neon-20, uranium-235 Nuclear Symbol

36. 3-3 Relative Atomic Mass • One atom has been chosen as a standard, masses of all other atoms expressed in relation to this standard • Atomic mass unit – exactly 1/12 the mass of a carbon-12 atom • Atomic mass of any other atom is determined by comparing it with the mass of carbon-12

37. 3-3 Average Atomic Mass • Most elements occur naturally as mixtures of isotopes • The percentage of each isotope in an element is constant – same no matter where sample comes from • Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element

38. 3-3 Calculating Average Atomic Mass Copper has 2 isotopes – copper-63 and copper-65 Copper-63: 69.17%, 62.929599 amu Copper-65: 30.83%, 64.927793 amu (.6917)(62.929599 amu) + (.3083)(64.927793 amu) = 63.55 amu

39. 3-3 Average Atomic Mass • Atomic masses that appear on the periodic table are AVERAGE atomic masses

40. 3-3 Relating Mass to Number of Atoms • THE MOLE! • The amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12 • The mole is a counting unit, like a dozen or a ream • 1 mole of something is 6.022 x 1023

41. 3-3 Avogadro’s Number • 12 g of carbon-12 contains exactly 6.0221367 x 1023 atoms. • Named Avogadro’s number after Amedeo Avogadro, Italian scientist

42. 3-3 Molar Mass • A mole is the amount of a substance that contains Avogadro’s number of particles • The mass of one mole of a pure substance is called the molar mass of that substance (g/mol)

43. 3-3 Molar Mass • To find the molar mass of an atom, find the atomic mass in amu on the periodic table and change the unit to g/mol • The molar mass is numerically equal to the atomic mass

44. 3-3 Gram/Mole Conversions • Molar mass can be used as a conversion factor • Can convert between mass and moles for any substance • Moles to Grams Multiply (by molar mass) • Grams to Moles Divide (by molar mass)

45. 3-3 Conversions with Avogadro’s Number • If amount in moles is known, can calculate number of particles (and vice versa) • Moles to Particles Multiply (by Avogadro’s number) • Particles to Moles Divide (by Avogadro’s number)

46. Mole Conversions