Chapter 4 Atoms

1. Chapter 4 Atoms Section 1 Development of Atomic Theory

2. BR. Who came up with the first theory of atoms? • Objectives- • 1. Give an example of how new scientific data can cause an existing scientific explanation to be supported, rejected or revised • 2. Evaluate selected theories based on supporting scientific evidence.

3. Objectives (cont.) • 3. Cite evidence that scientific investigations are conducted for many reasons. • 4. Identify scientific evidence that has caused modifications in previously accepted theories. • GLE’s:

4. Democritus • Over 2000 years ago • Democritus • Universe made of indivisible units called atoms. • Atomos- unable to be cut or divided • Did not have evidence to support theory

5. Dalton • 1808 John Dalton revised atomic theory • Atoms could not be divided • All atoms of a given element were exactly alike; and atoms of different elements could join to form compounds. • Based on experimental evidence

6. Dalton (cont.) • Law of definite proportions- a chemical compound always contains the same elements in exactly the same proportions by weight or mass. (supported Dalton’s theory) • Foundation of modern atomic theory • Could not explain all experimental evidence

7. Thomson • 1897 experiment suggested that atoms were not indivisible • He was experimenting with electricity • studying cathode rays not atoms • Cathode ray tube experiment suggested that cathode rays were made of negatively charged particles that came from inside the atoms

8. Thomson (cont.) • Revealed that atoms could be divided • Discovered electrons, negative particles • New model- electrons spread through-out the atom ( plum pudding model) Mass and positive charge evenly distributed Electrons scattered through out

9. Rutherford • Found Thomson’s model needing revising • Proposed that most of the mass of the atom was in the center • Conducted Gold-foil experiment where most particles passed straight through • Some particles were deflected • Some particles came straight back • Not what he expected

10. Rutherford (cont.) • Discovered the nucleus • Nucleus was very small • Electrons orbit the nucleus (like sun and planets) • Led to new model of the atom

11. Section 2- Structure of Atom • BR. Rutherford’s Gold foil experiment led to the discovery of what? • Objectives: • 1. Identify the 3 subatomic particles by location, charge, and relative mass • 2. Describe the results of loss or gain of electrons on charges of atoms

12. Objectives (cont.) • 3. Identify valence electrons in first 20 elements. GLE’S:

13. Atom • Three subatomic particles compared by mass, charge, and location in the atom. • Copy chart on page 119 • Nucleus- small, dense, center of atom • Atoms are Neutral.

14. Atom (cont.) • Nucleus is made of 1. protons- + charged particle 2. neutron- neutral or no charge particle Protons and neutrons are almost equal in size and mass.

15. Electrons • Move in a dense cloud (fan blades moving) outside the nucleus • Very tiny--- 1837 electrons = 1 neutron or proton • Negatively charged • Exact location cannot be determined. Speed and direction cannot be determined • Located by shading; shaded region is orbital: darker shading better chance to find

16. Protons • Each element has a unique number of protons • Elements are identified by the number of protons that they have in the nucleus of an atom • Atoms are neutral because they have the same number of p and e. They cancel each other out.

17. Ions • Atoms that have lost or gained electrons. • Lose electrons become positive. • Gain electrons become negative • If atoms lose or gain electrons they are not atoms, but IONS.

18. Atoms • Atoms are held together by an electric force + and – charges attract each other by an electric force This attraction is what holds the atom together just like the attractive force between solids and liquids.

19. Atoms and Elements • Atoms of different elements have unique structures. • Because atoms have different structures, they have different properties. • Atoms of the same element can vary in structure also. • Atoms of each element have the same number of protons, but different numbers of neutrons.

20. Atomic number • Atomic number equals the number of protons. • Since atoms are neutral, it also equals the number of electrons. • Neutral atom-- + = - Atomic numbers go from 1 to 116

21. Mass Number • Equals the total number of subatomic particles in the NUCLEUS of the atom. • Nucleus contains p and n. • Mass number is equal to p + n. • Neutrons vary so mass number can vary for the same element. • See figure 3 on page 121

22. Isotopes • Isotope has same atomic number but different number of neutrons. • Isotopes have same atomic number but different mass numbers. • Isotopes have the same number of protons but different number of neutrons

23. Isotopes (cont.) • Some isotopes are more common than others. • Radioisotopes- unstable isotopes that emit radiation and decay into other isotopes. They continue to decay until they reach a stable isotope. They decay at a fixed rate. (fraction of a second to millions of years)

24. Isotopes (cont.) • Since isotopes have the same number of protons and electrons, they have the same chemical properties. • Isotopes have different masses. • Isotopes of an element vary in mass because their numbers of neutrons differ.

25. Isotopes of Water • Water has 3 isotopes. Each isotope has 1 proton and 1 electron. • Protium- 1 proton and no neutrons (mass number of 1 • Deuterium-1 proton and 1 neutron (mass number of 2) • Tritium- 1 proton and 2 neutrons mass number of 3)

26. Calculating Neutrons • Isotopes are written 35 mass number ( p+n) • Cl symbol 17 atomic number (p) Mass number – atomic number = neutrons

27. Atomic mass • Atoms are expressed in unified atomic mass units because the mass is so small. • Unified atomic mass = equal to 1/12th of the mass of a carbon-12 atom. • Also called atomic mass unit

28. Atomic Mass • Average atomic mass for an element is a weighted average. • More common isotopes have more effect than less common isotopes of the element.

29. Mole • Mole- collection of a very large number of particles. 602,213,670,000,000,000,000,000 particles Written 6.02 x 1023 particles and is called Avogadro’s number Avogadro’s number= the number of atoms in 12 grams of carbon-12. (Popcorn kernels covering the US 310 miles tall)

30. Molar mass • Molar mass- the mass in grams of 1 mole of a substance 1 mole C12 = 12 g

31. Converting between Moles and Grams Amount xmolar mass of element= Mass(g Moles 1 mole of element 3 moles x32.07 g S = 96.21 g S S 1 mole S 96.27 g S x1 mole S= 3 mol S 32.21 g S Work problem 1 a-d on page 126

32. Compounds have Molar mass • Add all molar masses in compounds and then work the same way. H2O x(1.01 g H x 2) + 16 g O = 18.01 g/mol 1 mole H2O H2O 1 mol H20 = 18.02 g

33. Section 3Modern Atomic Theory • BR List the 3 subatomic particles and give their location, relative mass, and charge. • Objectives: • Describe the results of the loss or gain of electrons on the charges of atoms. • Identify valence electrons in first 20 elements. • Draw Bohr models of 1st 20 elements • GLE’S:

34. Modern Model of Atom • Electrons are found only in certain energy levels. NOT between levels • Location of electrons can not be predicted precisely • Bohr- electrons can be in only certain energy levels. • Bohr energy level related to electron’s path around the nucleus

35. Modern Model of Atom (cont.) • Electrons must gain energy to move to a higher energy level. • Electrons must lose energy to move to a lower energy level. • 1925 Bohr’s model revised.

36. Modern Atomic Theory (cont.) • Old out – not like sun and planets • New in – • Electrons behave more like waves on a vibrating string than like particles.

37. Energy levels and Electrons • Many energy levels for electron to occupy • The number of energy levels that are filled in an atom depends on the number of electrons. • Valence electrons- those electrons in outer most energy level • Valence electrons determine the chemical properties of the atom.

38. Energy levels • Maximum electrons in energy level • 1st = 2 • 2nd= 8 • 3rd= 18 • 4th= 32 • Must fill 1st and 2nd energy level before going to the 3rd energy level.

39. Energy levels (cont.) • There are 4 types of orbitals. • Orbitals are s, p, d, f. • Orbitals determine the number of electrons that each level can hold.

40. Electron Jumping • Electrons jump between energy levels when an atom gains or loses energy. • Lowest energy level called ground state. • Excited state-gains energy it moves to another level

41. How Electrons Move • Electrons gain energy by absorbing a photon and move to a higher energy level • Photon- particle of light; each have different energies • The electron may fall back to previous energy level when it releases a photon. • Photons determine which level the electron will • jump to.

42. Light- • Photons determine which level the electron will jump to. • Atoms absorb or emit light at certain wavelengths. • Energy of photon is related to the wavelength of the light. • High energy photons= short wavelengths • Low energy photons= long wavelengths

43. Atomic Fingerprint • Because of each element’s unique atomic structure, the wavelengths emitted depend on the particular element. • Each element emits its own characteristic color. Neon= red blue=copper sodium= yellow strontium= red orange = calcium green= barium