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 cm 3, 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 1.All matter is made of tiny, indivisible particles called atoms (in honor of Democritus’ idea) 2.All atoms of a given element have the same properties; atoms of different atoms have different properties 3.Compounds are formed by the chemical combinations of two or more different types of atoms 4.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 ParticlesParticleLocation Relative electrical charge Actual Mass (g) Electron In space surround- ing nucleus In space surround- ing nucleus 1- 1- 9.11 x 10 -28 Proton In nucleus 1+ 1+ 1.673 x 10 -24 Neutron In nucleus 0 1.675 x 10 -24 Relative Mass (amu) __1__ __1__ 1840 1840 1 1

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

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. Contains the symbol of the element, the mass number and the atomic number. X Mass number Atomic number

19. Examples 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 Mass number

20. Isotope Examples Isotope Mass number (A) Atomic number (Z) # protons # neutrons # electrons Carbon- 13 Cobalt - 58 Sodium- 23

21. Isotope Examples Isotope Mass number (A) Atomic number (Z) # protons # neutrons # electrons Carbon- 13 136676 Cobalt - 58 5827273127 Sodium- 23 2311111211

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 groups or families. Horizontal rows called periods.

27. Introduction to the Periodic Table IA IIAIIIAIVAVA VIA VII A Group number 1 2 3 5 4 6 7 Period number

29. 3 Main Categories of Elements 1.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 2.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. Stair Step Nonmetals Metals Metalloids

31. 3 Main Categories of Elements 3.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 Actinide series

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 2n 2

45. Energy levelmaximum number of (Shell) n electrons (2n 2 ) 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

51. d Orbitals 7 -

52. f-orbtals http://www.d.umn.edu/~pkiprof/ChemWebV2/AOs/ao4.html

53. Electron Configuration Electron configuration: Electron configuration: The arrangement of electrons in the extranuclear space (i.e. the empty space surrounding the nucleus). quantized The energy of the electrons in an atom is quantized, which means that an electron in an atom can have only certain allowed energies. These allowed energies correspond to specific regions in space surrounding the nucleus called energy levels or shells. Ground-state electron configuration: Ground-state electron configuration: The electron configuration of the lowest energy state of an atom.

54. Electron configurations Tell us the orbital location of an atom’s electrons Like an address

55. Electron Configuration Table 2.5 Distribution of Electrons in Shells

56. Assigning electrons to orbitals Orbital filling diagrams  use boxes or circles to represent orbitals See handout Rules for filling orbitals – Bottom up rule – atoms place their electrons in the lowest possible energy orbitals first – Each orbital can hold a maximum of 2 electrons, which must be spinning in opposite directions – For p, d and f orbitals, one electron in each orbital before pairing up

57. Electron Configuration Table 2.6 Distribution of Orbitals within Shells

58. Electron Configuration Figure 2.13 Energy levels for orbitals through the third shell.

59. Electron Configuration Electron configurations are governed by three rules: Rule 1: Rule 1: Orbitals fill in the order of increasing energy from lowest to highest. – Elements in the first, second, and third periods fill in the order 1s, 2s, 2p, 3s, and 3p.

60. Electron Configuration Rule 2: Rule 2: Each orbital can hold up to two electrons with spins paired in opposite directions. – With four electrons, the 1s and 2s orbitals are filled and are written 1s 2 2s 2. – With an additional six electrons, the three 2p orbitals are filled and are written either 2p x 2 2p y 2 2p z 2, or they may be written 2p 6.

61. Electron Configuration Orbitals have definite shapes and orientations in space

62. Electron Configuration Figure 2.14 The pairing of electron spins.

63. Electron Configuration Rule 3: Rule 3: When there is a set of orbitals of equal energy, each orbital becomes half filled before any of them becomes completely filled. – Example: – Example: After the 1s and 2s orbitals are filled, a 5th electron is put into the 2p x, a 6th into the 2p y, and a 7th into the 2p z. Only after each 2p orbital has one electron is a second added to any 2p orbital.

64. Electron Configuration Orbital box diagrams – A box represents an orbital. – An arrow represents an electron. – A pair of arrows with heads in opposite directions represents a pair of electrons with paired spins. Example Example: carbon (atomic number 6)

65. Electron Configuration Noble gas notation – The symbol of the noble gas immediately preceding the particular atom indicates the electron configuration of all filled shells Example: Example: carbon (atomic number 6)

66. Electron Configuration Valence shell: Valence shell: The outermost incomplete shell. Valence electron: Valence electron: An electron in the valence shell. Lewis dot structure: – The symbol of the element represents the nucleus and filled shells. – Dots represent valence electrons.

67. Electron Configuration

68. Table 2.9 Noble Gas Notation and Lewis dot structures for the Alkali Metals (Group 1A Elements)

69. Periodic Property As we have seen, the Periodic Table was constructed on the basis of trends (periodicity) in chemical properties. With an understanding of electron configuration, chemists realized that the periodicity of chemical properties could be understood in terms of periodicity in electron configuration. The Periodic Table worked because elements in the same column (group) have the same configuration in their outer shells. We look at two periodic properties: Atomic size and ionization energy.

70. Atomic Size The size of an atom is determined by the size of its outermost occupied orbital. Example: The size of a chlorine atom is determined by the size of its three 3p orbitals, the size of a carbon atom is determined by the size of if its three 2p orbitals.

71. Atomic Size Figure 2.16 Atomic radii of the main- group elements (in picometers).

72. Ionization Energy Ionization energy: Ionization energy: The energy required to remove the most loosely held electron from an atom in the gaseous state. – Example: When lithium loses one electron, it becomes a lithium ion; it still has three protons in its nucleus, but now only two electrons outside the nucleus, and therefore has a positive charge.

73. Ionization Energy Ionization energy is a periodic property: – In general, it increases across a row; valence electrons are in the same shell and subject to increasing attraction as the number of protons in the nucleus increases. – It increases going up a column; the valence electrons are in lower principle energy levels, which are closer to the nucleus and feel the nuclear charge more strongly.