1. 1 Modern Chemistry Chapter 3 Atoms: the building block of matter

2. 2 Chapter Vocabulary Law of conservation of mass Law of definite proportions Law of multiple proportions Atom Nuclear forces Atomic number Isotope Mass number nuclide Atomic mass unit Average atomic mass Mole Avogadro’s number Molar mass

3. Atomic Number Atoms are composed of identical protons, neutrons, and electrons –How then are atoms of one element different from another element? Elements are different because they contain different numbers of PROTONS The “atomic number” of an element is the number of protons in the nucleus # protons in an atom = # electrons

4. Chapter 3 Section 2 The Structure of the Atom pages 72-76 4 Composition of The Atomic Nucleus Nuclei contain protons and neutrons The nuclei is neutral because the number of protons equal the number of electrons Each element has a different number of protons in their nucleus –The number of protons determines the atom’s identity Nuclear forces hold protons & neutrons together

5. Chapter 3 Section 2 The Structure of the Atom pages 72-76 5 Properties of Subatomic Particles p. 76

6. Chapter 3 Section 2 The Structure of the Atom pages 72-76 6 Discovery of the Atomic Nucleus Relative size of the nucleus

7. Chapter 3 Section 3 Counting Atoms pages 77-87 7 Isotopes Atoms of the same element that have different masses Isotopes of hydrogen –Protium 1p + 0n 0 –Deuterium 1p + 1n 0 –Tritium 1p + 2n 0 Isotopes do not differ significantly in their chemical behavior

8. Chapter 3 Section 3 Counting Atoms pages 77-87 8 Mass Numbers Mass numbers = # of p + + # of n 0 of a specific isotope Examples –Protium 1p + + 0n 0 = 1 –Deuterium 1p + + 1n 0 = 2 –Tritium 1p + + 2n 0 = 3

9. Chapter 3 Section 3 Counting Atoms pages 77-87 9 Designating Isotopes Hyphen notation –name of element – mass number –Hydrogen – 3 Nuclear symbol mass number atomic number

10. Chapter 3 Section 3 Counting Atoms pages 77-87 10 Number of neutrons in an atom neutrons = mass number – atomic number How many p +, e - and n 0 are there in an atom of chlorine-37? 17 p + 17e- 20n 0 (37-17) Nuclide – a general term for a specific isotope of an element

11. Chapter 3 Section 1 Atoms: Ideas to Theory pages 67-71 11 Atoms: From Philosophical Idea to Scientific Theory

12. Chapter 3 Section 1 Atoms: Ideas to Theory pages 67-71 12 Foundation of Chemical Atomic Theory Law of Conservation of Mass –Mass is neither created or destroyed during ordinary chemical reactions or physical changes

13. Chapter 3 Section 1 Atoms: Ideas to Theory pages 67-71 13 Law of Conservation of Mass Image p. 69*

14. Chapter 3 Section 1 Atoms: Ideas to Theory pages 67-71 14 Law of Conservation of Mass Image p. 69*

15. The atom – Philosophical Idea to Scientific Theory 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

16. Dalton’s Atomic Theory (experiment based!) 3)Atoms of different elements combine in simple whole-number ratios to form chemical compounds 4)In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element. 1)All elements are composed of tiny indivisible particles called atoms 2)Atoms of the same element are identical. Atoms of any one element are different from those of any other element. John Dalton (1766 – 1844)

17. Sizing up the Atom  Elements are able to be subdivided into smaller and smaller particles – these are the atoms, and they still have properties of that element  If you could line up 100,000,000 copper atoms in a single file, they would be approximately 1 cm long  Despite their small size, individual atoms are observable with instruments such as scanning tunneling (electron) microscopes

18. Structure of the Nuclear Atom One change to Dalton’s atomic theory is that atoms are divisible into subatomic particles: –Electrons, protons, and neutrons are examples of these fundamental particles –There are many other types of particles, but we will study these three

19. Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron

20. Mass of the Electron 1916 – Robert Millikan determines the mass of the electron: 1/1840 the mass of a hydrogen atom; has one unit of negative charge The oil drop apparatus Mass of the electron is 9.11 x 10 -28 g

21. Conclusions from the Study of the Electron: a)Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. b)Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons c)Electrons have so little mass that atoms must contain other particles that account for most of the mass

22. Thomson’s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model. J. J. Thomson

23. Ernest Rutherford’s Gold Foil Experiment - 1911  Alpha particles are helium nuclei - The alpha particles were fired at a thin sheet of gold foil  Particles that hit on the detecting screen (film) are recorded

24. Chapter 3 Section 2 The Structure of the Atom pages 72-76 24 Gold Foil Experiment Image p. 75

25. Chapter 3 Section 3 Counting Atoms pages 77-87 25 Relative Atomic Mass One atom, carbon-12, is set as a standard All masses are expressed in relation to this standard 1 atomic mass unit = 1/12 the mass of a carbon-12 atom

26. Chapter 3 Section 3 Counting Atoms pages 77-87 26 Relative Atomic Mass Examples –Hydrogen – 1 = 1.007825 amu –Oxygen – 16 = 15.994915 amu –Magnesium – 24 = 23.985042 amu p + = 1.007276 amu, n 0 = 1.008665 amu, e - = 0.0005486 amu Relative mass and mass number are close in value but not the same

27. Chapter 3 Section 3 Counting Atoms pages 77-87 27 Average Atomic Mass The weighted average of the atomic masses of the naturally occurring isotopes of an element Example –Copper Cu-63:.6915 x 62.93 amu = 43.52 Cu-65:.3085 x 64.93 amu = 20.03 63.55 amu percent relative mass

28. Chapter 3 Section 3 Counting Atoms pages 77-87 28 The Mole An amount of a substance that contains as many particles as there are atoms in exactly 12 g carbon-12. Similar to a dozen or a pair or a gross 6.022 x 10 23 carbon-12 atoms = 12 grams of carbon-12 Avogadro’s number = 6.022 x 10 23 particles

29. Chapter 3 Section 3 Counting Atoms pages 77-87 29 Molar mass The mass of one mole of a pure substance Unit = g/mol

30. Chapter 3 Section 3 Counting Atoms pages 77-87 30 Gram-Mole Conversions The conversion factor for gram-mole conversion is molar mass. What is the mass, in grams, of 3.50 moles of Cu? –222 grams Cu OR g mol g

31. Chapter 3 Section 3 Counting Atoms pages 77-87 31 Conversions Image p. 84