1. Ch # 5 Atoms, Molecule, Formulas, Subatomic Particles.

2. Atoms • An atom is the smallest particle of an element that can exist and still have properties of the element.

3. Atomic Theory of Matter • All matter is made up of small particles called atoms.113 types. • Atoms of same element are similar to one another. • The relative number and arrangement of different types of atoms contained in a pure substance determined its identity.

4. Atomic Theory of Atoms. • These atoms rearrange, separate or unite to form a new substance. • Atoms can participate in or result from any chemical change, atoms cannot be destructed.

5. An atom is very Small

6. 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.

7. The Molecule • A group of two or more atoms that is strongly bound to each other.

8. The Molecule • Diatomic molecule: contains 2 atoms.H,N,O,F,Cl,Br,I, • Triatomic molecule: Contains 3 atoms. • P-4 atoms, S-8 atoms. • Homoatomic molecule: All atoms are of same kind. • Heteroatomic molecule: Two or more different kinds of atoms are present.

11. Molecular compounds • Compounds that have heteroatomic molecule as basic structural unit. • Ex: table sugar-sucrose. • Molecule is a limit of physical sub-division. Atom is a limit of chemical sub-division.

12. Molecular Compounds • Two kinds of atoms present, but, one kind of molecule. • Molecules have different properties than that of the atom.

13. Ionic compounds • Compounds that contain ions. Ex: NaCl

16. Natural and Synthetic compounds • The compounds could be naturally or artificially made.

17. Chemical Formulas • Made up of symbols. • The subscripts indicate the number of atoms. • Read the formula C6H12O6

18. Protons

19. 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.

20. Subatomic Particles. • Protons: +ve charge.-Goldstein. Charge: +1 • Electron:-ve charge.- Thomson. Smallest mass.Charge:-1 • Neutron: No charge.-Chadwick. • Opposite charges attract, like charges repel.

21. Nucleus: • Small, dense, positively charged center of an atom. It contains protons and neutrons. An atoms entire mass is concentrated to the center.

22. Nucleon • Any subatomic particle found in the nucleus of an atom. Ex; protons and neutrons. • Electrons revolve around the nucleus in a very large region. It is mostly empty space. • Volume occupied by electrons is called electron cloud. It is negatively charged.

23. An atom as a whole is neutral. Number of protons=Number of electrons in any atom. The +ve and-ve charges cancel each other out.

24. Diameter of nucleus= 10-15 m. Entire mass is concentrated in the center of the atom.

25. Protons and neutrons are made of Leptons, Mesons, Baryons.

28. 5.6

29. Atomic Numbers of the Elements

30. The atomic number of an element is equal to the number ofprotons in the nucleus of that element. • The atomic number of an atom determines which element the atom is.

31. Every atom with an atomic number of 1 is a hydrogen atom.Every hydrogen atom contains 1 proton in its nucleus.

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

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

34. Isotopic Notation

35. Isotopes of the Elements

36. Atoms of the same element can have different masses. • 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.

37. Isotopes of the Same Element Have Equal numbers of protons Different numbers of neutrons

38. Hydrogen has three isotopes 1 proton 0 neutrons 1 proton 1 neutron 1 proton 2 neutrons

39. Examples of Isotopes ElementProtonsElectronsNeutronsSymbol Hydrogen 1 1 0 11H Hydrogen 1 1 1 12H Hydrogen 1 1 2 13H Uranium 92 92 143 92235U Uranium 92 92 146 92238U Chlorine 17 17 18 1735Cl Chlorine 17 17 20 1737Cl

40. Atomic Mass

41. Using a mass spectrometer, the mass of the hydrogen atom was determined. • The mass of a single atom is too small to measure on a balance.

42. Using a mass spectrometer, the mass of one hydrogen atom was determined to be 1.673 x 10-24 g.

43. This number is very small. small small small small small small small small small small small small small small small small small small small small

44. Positive ions formed from sample. Electrical field at slits accelerates positive ions. Deflection of positive ions occurs at magnetic field. A Modern Mass Spectrometer From the intensity and positions of the lines on the mass spectrogram, the different isotopes and their relative amounts can be determined. A mass spectrogram is recorded. 5.8

45. The standard to which the masses of all other atoms are compared to was chosen to be the most abundant isotope of carbon.

46. 1 amu = 1.6606 x 10-24 g 1 amu is defined as exactly equal to the mass of a carbon-12 atom

47. Average atomic mass 1.00797 amu.

48. Average atomic mass 248.029 amu.

49. Average RelativeAtomic Mass

50. Most elements occur as mixtures of isotopes. • Isotopes of the same element have different masses. • The listed atomic mass of an element is the average relative mass of the isotopes of that element compared to the mass of carbon-12 (exactly 12.0000…amu).