Chapter 2 Atoms, Molecules, and Ions

1. Chapter 2Atoms, Molecules, and Ions

2. Introduction • Atoms • Composed of electrons, protons and neutrons • Molecules • Combinations of atoms • Ions • Charged particles

3. Greeks: Empedocles and Democritus • Suggested the concept of atoms but were not taken seriously or credited with an atomic theory

4. John Dalton: credited with the first atomic model

5. Figure 2.1 - John Dalton and Atomic Theory

6. Atomic Theory • An element is composed of tiny particles called atoms • All atoms of the same element have the same chemical properties 3. In an ordinary chemical reaction, atoms rearrange their bonds but atoms are not created or destroyed 4. Compounds are formed when two or more atoms of different element combine

7. Fundamental Laws of Matter

8. Law of Conservation of Mass Matter is conserved in chemical reactions This applies to all chemical reactions but DOES NOT include nuclear reactions

9. Law of Constant Composition Compound always contains the same elements in the same proportions by mass. Pure water has the same composition everywhere.

10. Law of Multiple Proportions • The masses of one element that combine with a fixed mass of the second element are in a ratio of small whole numbers. • Compare CO and CO2

11. Figure A – The Law of Multiple Proportions Two different oxides of chromium

12. Components of the Atom • Atomic theory raised more questions than it answered • Could atoms be broken down into smaller particles • 100 years after atomic theory was proposed, the answers were provided by experiment • Finding the Electrons: Protons: Neutrons:

13. J.J. Thomson • Discovered the electron

14. Figure 2.2 – J.J. Thomson and Ernest Rutherford

15. Figure 2.3 – Cathode Ray Apparatus

16. Electrons • First evidence for subatomic particles • J.J. Thomson in 1897 • Rays emitted were called cathode rays • Rays are composed of negatively charged particles called electrons • Electrons carry unit negative charge (-1) and have a very small mass (1/2000 the lightest atomic mass)

17. J.J. Thomson’s Model • Every atom has at least one electron • Atoms are known that have one hundred or more electrons • There is one electron for each positive charge in an atom • Electrical neutrality is maintained

18. Ernest Rutherford: Discovered the nucleus of the atom

19. Gold Foil Experiment: • Bombardment of gold foil with αparticles (helium atoms minus their electrons) • Expected to see the particles pass through the foil • Found that some of the alpha particles were deflected by the foil • Led to the discovery of a region of heavy mass at the center of the atom = nucleus

20. Figure 2.4 – Rutherford Backscattering

21. Nuclear Particles 1. Protons • Mass nearly equal to the H atom • Positive charge 2. Neutrons • Mass slightly greater than that of the proton • No charge

22. Atomic Mass • The average mass of all of the isotopes of an element accounting for their relative abundances

23. Table 2.1 – Subatomic Particles

24. Terminology • Atomic number, Z • Number of protons in the atom • Mass number, A • Number of protons plus number of neutrons • Mass # = p+ + n0

25. A is the mass number Z is the atomic number X is the chemical symbol Nuclear symbolism

26. Isotopes • Isotopes are two atoms of the same element • Same atomic number but differ in number of neutrons • Different mass numbers • Mass # = p+ + n0

27. Example 2.1

28. Radioactivity • Radioactive isotopes are unstable (Radioactive decay is not a chemical process) 1. These isotopes decay over time 2. Emit other particles and are transformed into other elements • Particles emitted 1. Beta (β) particles: High speed electrons 2. Alpha (α) particles: helium nuclei 3. Gamma (γ) rays: high energy light

29. Nuclear Stability • depends on the neutron/proton ratio • For light elements, n/p is approximately 1/1 • For heavier elements, n/p is approximately 1.4/1

30. Figure 2.5 – The Nuclear Belt of Stability

31. 2.3 Introduction to the Periodic Table • Dmitri Mendeleev: 1836-1907 • Arranged elements by chemical properties • Left space for elements unknown at the time • Predicted detailed properties for several undiscovered elements: • Sc, Ga, Ge • By 1886, all these elements had been discovered, and with properties similar to those he predicted

32. Mendeleev’sP.T.

33. Introduction to the Periodic Table

34. Modern Periodic Table • Period – a horizontal row on the periodic table • Group – a vertical column on the periodic table • Blocks – sections of elements with common properties • Families – another name for group; emphasizes the similarity in properties within a group

35. Blocks in the Periodic Table • Main group elements • 1-2, 13-18 OR roman numeral +A groups • Transition Metals • 3-12 OR non roman numeral groups • Inner Transition/Rare Earth elements • Bottom double rows

36. Families with Common Names (label on PT) • Alkali Metals, Group 1(I) • Alkaline Earth Metals, Group 2 (II) • Halogens, Group 17 (VII) • Noble Gases, Group 18 (VIII)

37. A Look at the Sulfur Group • Sulfur (nonmetal), antimony (metalloid) and silver (metal)

38. Example 2.3

39. 2.4 Molecules and Ions • Molecule: Two or more atoms chemically combined 1. Atoms involved are often nonmetals 2. Covalent bonds are strong forces that hold the atoms together • Molecular formulas: • Number of each atom is indicated by a subscript • Examples • Water, H2O • Ammonia, NH3

40. Structural Formulas • Structural formulas: a formulas that shows the bonding patterns within the molecule

41. Ions • A charged particle that is the result of the loss or gain of electrons • Cation – a positive ion (loss) • Anion – a negative ion (gain) • Examples: • Na → Na+ + e- • O + 2e-→ O2-

42. Ionic Compounds • Compounds formed from the electrostatic attraction of oppositely charged particles • Sodium chloride (NaCl): Sodium cations and chloride anions associate into a continuous network

43. Forces: • Ionic compounds are held together by strong forces • Compounds are usually solids at room temperature • High melting points • often water-soluble

44. Solutions: • When an ionic compound dissolves in water, the ions are released from each other • conductivity – the ions in a solution support the transmission of an electric current • Strong electrolytes – solutions that are very good conductors • Weak electrolytes – solutions that are poor conductors • Nonelectrolytes – solutions that do NOT conduct

45. Figure 2.12 – Electrical Conductivity

46. Formulas for Ionic Compounds • Charge balance • Each positive charge must have a negative charge to balance it • Calcium chloride, CaCl2 • Ca2+ • Two Cl- ions are required for charge balance

47. Transition Metals • Polyvalent – exhibit multiple positive charges depending on conditions • Iron forms Fe2+ and Fe3+ • Lead forms Pb2+ and Pb4+

48. Polyatomic Ions • Groups of atoms may carry a charge; these are the polyatomic ions • OH- • NH4+

49. Noble Gas Connections • Atoms that are close to a noble gas (group 18 or VIII) form ions that contain the same number of electrons as the neighboring noble gas atom • +1, +2, +3 skip -3, -2, -1 Noble Gases

50. Example 2.5