Chapter 2

1. Chapter 2 Atoms

2. What is Chemistry? • The study of matter and its properties and transformations What is Matter? • Anything that has mass and volume • Mass = the amount of a substance, measured in grams, g • Volume = the space occupied by a substance, measured in cm3, mL (milliliters) or L (Liters)

3. Brief History of Chemistry • Ideas about matter date back to ancient Greece  2000 years ago • 2 schools of thought Democritus – all matter made of tiny, indivisible particles called atoms (from Greek word atomos = uncuttable) Aristotle – matter is continuous, it is infinitely divisible • Aristotle’s ideas dominated for almost 2000 yrs

4. So, who was right? • Today we know that Democritus was right • Atoms are the basic building blocks of matter • we will discuss evidence for the existence of atoms later

5. Why do we believe in atoms? • First atomic theory based on scientific evidence proposed in 1808 by English chemist John Dalton (1766 -1844) • Theory based on three scientific laws discovered in late 1700s, early1800s • Law of Conservation of Mass (Antoine Lavoisier (1743-1794)) • Matter cannot be created or destroyed in an ordinary chemical reaction

6. Law of Constant Composition (Joseph Proust (1754-1826)) • no matter where you find a specific compound, it is always made up of the same proportion of elements by mass  elements combine to form compounds in fixed proportions

7. Law of Multiple Proportions (John Dalton) • elements always combine to make compounds in whole number ratios or multiples of whole number ratios, never in fractions (Not mentioned in book)

8. Postulates of Dalton’s Atomic Theory • All matter is made of tiny, indivisible particles called atoms (in honor of Democritus’ idea) • All atoms of a given element have the same properties; atoms of different atoms have different properties • Compounds are formed by the chemical combinations of two or more different types of atoms • During chemical reactions, atomic arrays are just rearranged into new combinations

9. Dalton’s Model of the Atom • Atoms are solid, indivisible spheres, like billiard balls • His model was referred to as the “Billiard Ball” model • Dalton’s model of the atom was to endure for almost 100 years, until the discovery of radioactivity and the first subatomic particle (the electron) in the late 1890s.

10. What are atoms made of? • First subatomic particle, the electron, discovered by English Physicist J.J. Thomson in 1897. • The proton was discovered in 1919 by Ernest Rutherford • The last subatomic particle to be identified was the Neutron in 1932 by James Chadwick.

13. Properties of Subatomic Particles

14. Terminology for Atomic Structure Atomic number (Z) – the number of protons in the nucleus of an atom, also the number of electrons as atoms are electrically neutral Mass number (A) – the number of protons and neutrons in the nucleus of an atom • Number of neutrons in the nucleus : #neutrons = mass no. – atomic no.

15. Subatomic Particles Determining the number of protons and electrons in an atom from the periodic table Atomic number = # protons = 6 = # electrons = 6 12.01 C 6 6 C Carbonn Atomic number Symbol

16. Subatomic Particles Determining the number of neutrons in an atom: Mass # - Atomic # • Must be given the mass number! • Mass number is not the same as the atomic mass • e.g. Sodium with a mass number of 23 Na atomic # = 11, 11 protons, 11 e- neutrons = 23 – 11 = 12 neutrons

17. Discovery Of Isotopes • After neutrons discovered, it was found that not all atoms of the same element were the same (as Dalton had said) • Almost every element has examples of atoms that have the same number of protons, but different numbers of neutrons Isotopes = Atoms that have the same number of protons (atomic number), but different numbers of neutrons (different mass numbers)

18. Nuclear Notation • Contains the symbol of the element, the mass number and the atomic number. Mass number X Atomic number

19. Examples Mass number C Number of protons = 6 Number of electrons = 6 Number of neutrons = 12 – 6 = 6 Mass number can also be used at the end of the element’s name e.g. carbon-12 12 6 Atomic number

20. Isotope Examples

21. Isotope Examples

22. Introduction to the Periodic Table

23. Dmitri Mendeleev (1834-1907) • When Mendeleev arr- anged the elements by increasing atomic weight, he noticed a periodic repetition in atomic properties (e.g. density, melting point)

24. Because the properties of the atoms were repeated periodically, the table he created was called periodic table

25. Mendeleev’s 1872 Table

26. Modern Periodic Table • After Moseley discovered atomic number, elements were rearranged from increasing atomic weight to increasing atomic number • Vertical Columns called groupsor families. • Horizontal rows called periods.

27. Introduction to the Periodic Table Period number Group number IA VIIA 1 IIA IIIA IVA VA VIA VIIA 2 3 4 5 6 7

29. 3 Main Categories of Elements • Metals – Shiny, good conductors of electricity and heat, tend to have 3or less valence e-, malleable, ductile, located to the left of the stair step on the periodic table • Nonmetals – dull, brittle, poor (some non) conductors of electricity and heat, have 4 or more valence e-, located to the right of the stair step on the periodic table

30. Metalloids Nonmetals StairStep Metals

31. 3 Main Categories of Elements • Metalloids – located along the stair step on the periodic table, have properties of both metals and nonmetals e.g. Silicon – is shiny, brittle, semiconductor of electricity

32. Introduction to the Periodic Table Alkali metals

33. Introduction to the Periodic Table Alkaline earth metals

34. Introduction to the Periodic Table Halogens

35. Introduction to the Periodic Table Noble gases

36. Introduction to the Periodic Table Transition metals

37. Introduction to the Periodic Table Lanthanide series Actinideseries

38. Introduction to the Periodic Table Metalloids

39. Introduction to the Periodic Table Main Group Elements

40. Periodicity • Trend within a group of elements or across a period of elements in the periodic table

41. How are the electrons in an atom arranged? • Atom is mostly empty space with central dense core called nucleus • Electrons are located at a distance away and have to be constantly moving to avoid being pulled into the positively charged nucleus • Because e- are moving, they possess kinetic energy • In 1913, Niels Bohr discovered

42. Niels Bohr 1913 • Discovered that only certain values are possible for the energy of the hydrogen electron • The energy of the electron is quantized  only certain values are allowed

43. Quantized Energy Levels • The energy levels of all atoms are quantized.

44. Electrons are confined to specific regions of space, called principal energy levels or shells • These energy levels or shells radiate away from the nucleus and given whole integer numbers of 1, 2, 3, 4, etc • Each energy level can accommodate only a certain number of electrons, given by the formula 2n2

45. Energy level maximum number of (Shell) n electrons (2n2) 1 2 2 8 3 18 4 32

46. Energy levels are further divided into sublevels or subshells • Sublevels are designated by the letters s, p, d and f n = 1  1 sublevel = 1s n = 2  2 sublevels = 2s and 2p n= 3  3 sublevels = 3s, 3p and 3d n = 4  4 sublevels = 4s, 4p, 4d, and 4f

47. Within these sublevels, electrons are grouped in orbitals Orbital = most probable region in space of finding an electron • According to quantum theory, there is a limit to what we can know about the electron • Therefore, we can only discuss its location in terms of probability. • Orbitals are probability maps that have definite shapes and orientations in space • Each orbital can hold a maximum of 2 electrons

48. Sublevel designation s, p, d and f also designates the shape of the electron orbital S orbitals = spherical an shape 1s 2s 3s

49. p orbitals • p orbitals are dumbbell shaped • There are three p orbital shapes

50. The s and p types of sublevel