Early Chemistry. Early Chemists only believed in 1 element: Dirt. Later Chemists believed in 4 elements:

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2 Early Chemistry Early Chemists only believed in 1 element: Dirt Later Chemists believed in 4 elements: Air Earth Fire Water Various combinations of these produced various compounds

3 Atomic Structure All matter is composed of atoms. Understanding the structure of atoms is critical to understanding the properties of matter. properties of solid materials depend on the geometrical atomic arrangements, and the interactions between constituent atoms.

4 history of the atom 460 BC Democritus developed the idea of atoms! he pounded up materials in his pestle and mortar until he had reduced them to smaller and smaller particles which he called ATOMA (greek for indivisible) It took ~2400 years from when it was conceived to the time experimental evidence prove of the atom existence.

5 history of the atom 1808 John Dalton suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them ATOMS

6 history of the atom 1898 Joseph John Thompson / Cambridge found that atoms could sometimes eject a far smaller negative particle which he called ELECTRON 1906 Nobel prize in Physics

7 History of the atom 1910 Ernest Rutherford / Cambridge student of Thompson 1908 Nobel prize in Chemistry proposed a more detailed model with a central nucleus: positive charge was all in a central nucleus. With this holding the electrons in place by electrical attraction

8 History of the atom 1913 Niels Bohr / Danish / a football fanatic 1922 Nobel prize in physics studied under Rutherford at the Victoria University in Manchester. Bohr refined Rutherford's idea by adding that the electrons were in orbits. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons.

9 Bohr s atom Rutherford s model predicted a rainbow of colors rather than discrete lines obtained from an atomic line spectra. To explain the line spectra, Bohr proposed that electrons of specific energy moved in circular orbits around the nucleus and could not exist between these orbits.

10 Atomic Structure Atoms are composed of protons positively charged particles neutrons neutral particles nucleus electrons negatively charged particles in orbitals surrounding the nucleus.

11 Atomic Structure Every different atom has a characteristic number of protons in the nucleus. atomic number (Z) = number of protons For an electrically neutral atom, atomic number = number of electrons. Atoms with the same atomic number have the same chemical properties and belong to the same element.

12 Atomic Structure Z ranges from 1 for hydrogen to 92 for uranium (the highest for the naturally occuring elements).

13 Atomic Structure The atomic mass (A) of a specific atom: the sum of the number of protons and neutrons within the nucleus. mass number: A = Z + N number of protons is the same for all atoms of a given element, number of neutrons (N) may be variable.

14 Atomic Structure The number of protons in the nucleus of the atom is equal to the atomic number (Z). The number of electrons in a neutral atom is equal to the number of protons. The mass number of the atom (M) is equal to the sum of the number of protons and neutrons in the nucleus. The number of neutrons is equal to the difference between the mass number of the atom (M) and the atomic number (Z).

15 Atomic Structure/isotopes atoms of some elements have two or more different atomic masses, called isotopes.

16 Atomic Structure Atomic weight: Weighted average of the atomic masses of the atom s naturally occurring isotopes. Boron consists of the isotopes: 19.7% B-10 (5p+5n) and 80.3% B-11 (5p+6n). atomic weight for Boron, B = (19.7 x 10)+(80.3 x 11)]/100= 10.8 amu Bromine isotopes: 50.5% Br-79 and 49.5% Br-81. atomic weight for Bromine, Br = [(50.5 x 79)+(49.5 x 81)]/100= 80.0 amu

17 Atomic Structure In one mole of a substance there are x10 23 (Avogadro s number) atoms or molecules. Atomic weight = weight of x atoms For example, the atomic weight of iron is amu/atom, or g/mol. 1 amu/atom = 1g/mol = 1 dalton Atomic mass unit (amu): 1 12 of the atomic mass of carbon atomic mass of C: amu / of H:1.008 amu

18 subatomic particles particle Mass (g) Charge (C/eV) Electron (e-) 9.11x x Proton (p) 1.67x x Neutron (n) 1.67x Proton is 1837 times heavier than an electron. Neutron is 1842 times heavier than an electron. Electron is much lighter with respect to the protons and neutrons

19 Atomic Structure HELIUM ATOM proton nucleus # electrons = # protons - + N N + - Shell electron neutron ATOMIC MASS NUMBER = number of protons + number of neutrons ATOMIC NUMBER = number of protons

20 Lithium Protons Electrons Neutrons three electrons three protons four neutrons.

21 Beryllium Protons Electrons Neutrons four electrons four protons five neutrons.

22 Boron Protons Electrons Neutrons five electrons five protons six neutrons.

23 Carbon Protons Electrons Neutrons six electrons six protons six neutrons.

24 Nitrogen Protons Electrons Neutrons seven electrons seven protons seven neutrons.

25 Oxygen Protons Electrons Neutrons eight electrons eight protons eight neutrons

26 Fluorine Protons Electrons Neutrons nine electrons nine protons ten neutrons.

27 Neon Protons Electrons Neutrons ten electrons ten protons ten neutrons

28 Sodium Protons Electrons Neutrons eleven electrons eleven protons twelve neutrons

29 Atomic structure How many protons, neutrons and electrons?

30 Atomic structure How many protons, neutrons and electrons?

31 Atomic structure How many protons, neutrons and electrons?

32 Atomic structure How many protons, neutrons and electrons?

33 Atomic structure How many protons, neutrons and electrons?

34 Atomic structure The charge and mass number of an electron are: a) charge = 0, Mass number = 1 b) charge = -1, Mass number = 0 c) charge = +1, Mass number = 1 d) charge = +1, Mass number = 0 The charge and mass number of a neutron are? a) charge = +1, Mass number = 1 b) charge = 0, Mass number = 1 c) charge = +1, Mass number = 0 d) charge = -1, Mass number = 0

35 Atomic structure Which of the following has 25 protons and 31 neutrons? a) 56 Mn b) 56 Ga c) 25 Ga d) 31 Mn e) 56 Ba

36 Atomic structure Why does chlorine have an atomic mass of 35.5, which is not a whole number? a) Chlorine contains an extra electron which makes it weigh more than 35. b) Chlorine contains 17 protons and 18.5 neutrons c) Chlorine normally exists in an excited state, and so it weighs more than 35. d) The chlorine was not pure when its atomic mass was measured. e) Chlorine, as found in nature, contains a mixture of the isotopes 35 Cl and 37 Cl, in such proportions as to give an average atomic mass of 35.5

37 Atomic structure The two main parts of an atom are? a) nucleus and electron energy levels b) nucleons and protons c) oxidation number and valence d) protons and neutrons e) protons and electrons

38 Atomic structure The nucleus of the element having atomic number 25 and atomic weight 55 will contain? a) 25 protons and 30 neutrons b) 30 protons and 25 neutrons c) 55 protons d) 55 neutrons

39 Atomic structure A beryllium atom has 4 protons, 5 neutrons, and 4 electrons. What is the mass number of this atom? a) 4 b) 5 c) 8 d) 9 e) 13

40 Atomic structure The smallest particle into which an element can be divided and still have the properties of that element a) nucleus b) electron c) atom d) neutron How would you describe the nucleus? a) dense, positively charged b) mostly empty space, positively charged c) tiny, negatively charged d) dense, negatively charged

41 Atomic structure Where are electrons likely to be found? a) in the nucleus b) in electron clouds c) mixed throughout an atom d) in definite paths Every atom of a given element has the same number of a) protons b) neutrons c) electrons d) isotopes

42 Atomic structure What is the meaning of the word atom? a) dividable b) invisible c) hard particles d) not able to be divided Which statement is true about isotopes of the same element? a) They have the same number of protons b) They have the same number of neutrons c) They have a different atomic number d) They have the same mass

43 Atomic structure Which has the least mass in an atom? a) nucleus b) proton c) neutron d) electron If an isotope of uranium, uranium-235, has 92 protons, how many protons does the isotope uranium-238 have? a) 92 b) 95 c) 143 d) 146

44 Atomic structure What is the atomic mass number of a Ba atom? a) 56 b) 81 c) 137 d) 25

45 Atomic structure A carbon atom with 6 protons, 6 electrons, and 6 neutrons would have a mass number of a) 6 b) 12 c) 15 d) 18 The number at the top is the a) atomic number b) element name c) atomic mass d) chemical symbol

46 Atomic structure How many electrons does a neutral Cl atom contain? a)16 b)17 c)18 d)19 What is the difference between atomic mass and atomic weight? Atomic mass is the mass of a single atom or an individual isotope. The atomic weight is the average mass of all naturally occurring isotopes of an element.

47 Atomic Structure Neutral atoms have the same number of protons and electrons. Ions are charged atoms. cations have more protons than electrons and are positively charged anions have more electrons than protons and are negatively charged

48 Atomic Structure If a neutral atom looses one or more electrons it becomes a cation. Na 11 protons e- + Na + 11 protons 11 electrons 10 electrons If a neutral atom gains one or more electrons it becomes an anion. Cl 17 protons + e- Cl - 17 protons 17 electrons 18 electrons

49 Bohr Atomic model electrons are assumed to revolve around the atomic nucleus in discrete orbitals, and the position of any particular electron is more or less well defined in terms of its orbital. Electrons are permitted to have only specific values of energy.

50 Bohr Atomic model

51 excitation vs relaxation An electron may change energy by making a quantum jump either to an higher energy (with absorption of energy) or to a lower energy (with emission of energy). relaxation excitation

52 Quantum Mechanics Unfortunately, extremely small particles (electrons) do not follow the laws of classical (Newtonian) physics. The new physics that mathematically treats small particles is called Quantum Mechanics.

53 electron distribution wave-mechanical model an electron is no longer treated as a particle moving in a discrete orbital; electron is considered to exhibit both wavelike and particle-like characteristics. The position of an electron is described by a probability distribution // electron cloud.

54 Quantum Mechanics Wave behavior is described with the wave function ψ, incorporating the wave and particle features of electrons (Erwin Schrödinger) The probability of finding an electron in a certain area of space is proportional to ψ 2 electron density. Austrian; 1933 Nobel prize in physics

55 Quantum Mechanics

56 Heisenberg s uncertainty principle more precisely the position of some particle is determined, the less precisely its momentum can be known A macroscale analogy High Shutter Speed Can judge location, but not speed. Low Shutter Speed Can judge speed, But not location

57 Heisenberg s uncertainty principle we cannot precisely measure the momentum and the position of an electron at the same time. As the momentum of the electron is more and more certain, the position of the electron becomes less and less certain, and vice versa. n = 2.5 cannot exist as a principal quantum number. There must be an integral number of wavelengths (n) in order for an electron to maintain a standing wave. If there were to be partial waves, the whole and partial waves would cancel each other out and the particle would not move.

58 Quantum Mechanics The Schrödinger equation specifies possible energy states an electron can occupy. The energy states and wave functions are characterized by a set of quantum numbers. Instead of orbits in the Bohr model, quantum numbers and wave functions describe atomic orbitals in quantum mechanics.

59 quantum numbers every electron in an atom is characterized by four quantum numbers. There are three quantum numbers necessary to describe an atomic orbital. The principal quantum number (n) designates size The angular moment quantum number (l) describes shape The magnetic quantum number (m l ) specifies orientation

60 Principal Quantum Number (n) n designates the size of the orbital. Larger values of n correspond to larger orbitals. The allowed values of n are integers: 1, 2, 3 and so forth. A collection of orbitals with the same value of n is frequently called a shell. n K L M N O P

61 Angular moment Quantum Number (l) l signifies the subshell l describes the shape of the orbital. l values range from 0 to n 1 Example: If n = 2, l can be 0 or 1. n l subshell 0 0,1 0,1,2 0,1,2, s s,p s,p,d s,p,d,f s,p g s,p h energy state

62 Magnetic Quantum Number (m l ) describes the orientation of the orbital in space. m l are integers that depend on l: l, 0, +l m l identifies # of energy states for each subshell For an s subshell: a single energy state For p, d, and f subshells: 3, 5, and 7 energy states

63 Number of available electron states for initial shells and subshells Principal Quantum No: n Shell 1 K 2 L 3 M 4 N 5 O 6 P Subshell No. of energy States: l m l s /0 1 / 0 2 s / 0 1 / 0 2 p / 1 3 / -1,0,+1 6 s / 0 1 / 0 2 p / 1 3 / -1,0,+1 6 d / 2 5 / -2,-1,0,+1,+2 10 s / 0 1 / 0 2 p / 1 3 / -1,0,+1 6 d / 2 5 / -2,-1,0,+1,+2 10 f / 3 7 / -3,-2,-1,0,+1,+2,+3 14 s / 0 1 / 0 2 p / 1 3 / -1,0,+1 6 d / 2 5 / -2,-1,0,+1,+2 10 f / 3 7 / -3,-2,-1,0,1,2,3 14 g / 4 9 / -4,-3,-2,-1,0,+1,+2,+3,+4 18 s / 0 1 / 0 2 p / 1 3 / -1,0,+1 6 d / 2 5 / -2,-1,0,+1,+2 10 f / 3 7 / -3,-2,-1,0,1,2,3 14 g / 4 9 / -4,-3,-2,-1,0,1,2,3,4 18 h / 5 11 / -5,-4,-3,-2,-1,0,+1,+2,+3,+4,+5 22 Number of Electrons Per Subshell Per Shell

64 Atomic orbitals An s subshell has one orbital which is spherically shaped. If you were to measure where the electron was within an s subshell many, many times and plot the results on a graph you would get something like this.

65 Atomic Orbitals p orbitals: a dumbbell shape with electrons on either side of the nucleus in tear drop shaped lobes Three orientations: l = 1 (as required for a p orbital) m l = 1, 0, +1

66 Atomic Orbitals The d orbitals: Five orientations: l = 2 (as required for a d orbital) m l = 2, 1, 0, +1, +2

67 Atomic orbitals d-orbitals are followed by the seven f-orbitals. 7 orientations: l = 3 (as required for a d orbital) m l = -3, 2, 1, 0, +1, +2, +3

68 Quantum Numbers To summarize quantum numbers: principal (n) size angular (l) shape magnetic (m l ) orientation Required to describe an atomic orbital principal (n = 2) 2p x related to the magnetic quantum number (m l ) angular momentum (l = 1) electron spin (m s ) direction of spin Required to describe an electron in an atomic orbital

69 Electron Spin Quantum Number-ms used to specify an electron s spin. There are two possible directions of spin. Allowed values of m s are +½ and ½.

70 Pauli exclusion principle No Two Electrons in an Atom Can Have the Same Four Quantum Numbers; the same values for n, l, m l, and m s. Although the first three quantum numbers identify a specific orbital and may have the same values, the fourth is significant and must have opposite spins. a set of quantum numbers is specific to a certain electron.

71 Quantum numbers / Q An electron with n = 2, l = 1, ml = 1, and ms = +1/2 is found in the same atom as a second electron with n = 2, l = 1, ml = 1. What is the spin quantum number for the second electron? Since the first three quantum numbers are identical for these two electrons, we know that they are in the same orbital. As a result, the spin quantum number for the second electron cannot be the same as the spin quantum number for the first electron. This means that the spin quantum number for the second electron must be ms = 1/2.

72 Quantum numbers / Q An electron with n = 5, l = 4, m l = 3, and m s = 1/2 is found in the same atom as a second electron with n = 5, l = 4, m l = 3. m s =? Since the first three quantum numbers are identical for these two electrons, we know that they are in the same orbital. As a result, the spin quantum number for the second electron cannot be the same as the spin quantum number for the first electron. This means that the spin quantum number for the second electron must be ms = +1/2.

73 Quantum numbers / Q Can an electron with n = 1, l = 0, ml = 0, and ms = +1/2 exist in the same atom with a 2nd electron with n = 2, l = 0, ml = 0, and ms = +1/2? Since these two electrons are in different orbitals, they occupy different regions of space within the atom. As a result, their spin quantum numbers can be the same, and thus these two electrons can exist in the same atom.

74 Atomic structure Maximum number of electrons in a subshell with l = 3 and n = 4 is a) 10 b) 12 c) 14 d) 16 e) 18 Principal Quantum No: n 4 Subshell l No. of energy States: m l Number of Electrons Per Subshell s / 0 1 / 0 2 p / 1 3 / -1,0,+1 6 d / 2 5 / -2,-1,0,+1,+2 10 f / 3 7 / -3,-2,-1,0,+1,+2,+3 14

75 Atomic structure The lowest principal quantum number for an electron is? a) 0 b) 1 c) 2 d) 3 e) 4

76 Atomic structure Which sublevel can by occupied by a maximum of 10 electrons? a) s b) p c) d d) f Subshell l No. of energy States: m l Number of Electrons Per Subshell s / 0 1 / 0 2 p / 1 3 / -1,0,+1 6 d / 2 5 / -2,-1,0,+1,+2 10 f / 3 7 / -3,-2,-1,0,+1,+2,+3 14

77 Atomic structure The K, L and M shells of an atom are full. Its atomic number is. a) 18 b) 20 c) 10 d) 12 Principal Quantum No: n Shell 1 K 2 L 3 M Subshell No. of energy States: l m l s /0 1 / 0 2 s / 0 1 / 0 2 p / 1 3 / -1,0,+1 6 s / 0 1 / 0 2 p / 1 3 / -1,0,+1 6 d / 2 5 / -2,-1,0,+1,+2 10 Number of Electrons Per Subshell Per Shell

78 Atomic structure If n=3, and l=2, then what are the possible values of m l? Since ml must range from l to +l, then ml can be: -2, -1, 0, 1, or 2.

79 Atomic structure State whether an electron can be described by each of the following sets of quantum number. If a set is not possible, state why not. a) n = 2, l = 1, m l = -1 b) n = 1, l = 1, m l = +1 c) n = 4, l = 3, m l = +3 d) n = 3, l = 1, m l = -3

80 Atomic structure Replace the question marks by suitable responses in the following quantum number assignments. a) n = 3, l = 1, m l =? b) n = 4, l =?, m l = -2 c) n =?, l = 3, m l =?

81 Atomic structure Replace the question marks by suitable responses in the following quantum number assignments. a) n = 3, l = 1, m l = -1,0,1 b) n = 4, l = 2, m l = -2 c) n = 4, l = 3, m l = -3,-2,-1,0,1,2,3 Principal Quantum No: n Subshell l No. of energy States: m l 1 s /0 0 2 s / 0 0 p / 1-1,0,+1 s / p / 1-1,0,+1 d / 2-2,-1,0,+1,+2 4 f / 3-3,-2,-1,0,+1,+2,+3

82 Atomic structure / Q Provide the three quantum numbers describing each of the three p orbitals in the 2p subshell. n l m l 2p x p y 2p z Principal Quantum No: n Subshell l No. of energy States: m l 2 s / 0 0 p / 1-1,0,+1

83 Atomic structure For n = 1, determine the possible values of l. For each value of l, assign the appropriate letter designation & determine the possible values of m l. n = 1 l = 0 s m l = 0 Principal Quantum No: n Subshell l No. of energy states: m l 1 s /0 0

84 Atomic structure How many orbitals in shell n = 1? 1st Shell has only the s orbital! How many electrons possible? S orbital can hold only 2 electrons!

85 Atomic structure For n = 2, determine the possible values of l. For each value of l, assign the appropriate letter designation & determine the possible values of m l. Principal Quantum No: n Shell 2 L Subshell How many orbitals in shell n = 2? How many electrons possible? Principal Quantum No: n Shell 2 L l No. of energy States: m l s / 0 0 p / 1-1, 0, +1 Subshell l s / 0 2 p / 1 6 Number of Electrons Per Subshell Per Shell 8

86 Atomic structure For n = 3, determine the possible values of l. For each value of l, assign the appropriate letter designation & determine the possible values of m. l n Shell l m l 3 M s / 0 1 / 0 p / 1 3 / -1,0,+1 d / 2 5 / -2,-1,0,+1,+2 How many orbitals in shell n = 3? How many electrons possible? n Shell l Number of Electrons Per Subshell Per Shell 3 M s / 0 2 p / 1 6 d /

87 Atomic structure For n = 4, determine the possible values of l. For each value of l, assign the appropriate letter designation & determine the possible values of m l. Principal Quantum No: n Shell 4 N Subshell l No. of energy States: m l s / 0 1 / 0 p / 1 3 / -1,0,+1 d / 2 5 / -2,-1,0,+1,+2 f / 3 7 / -3,-2,-1,0,+1,+2,+3

88 Atomic structure Provide the four quantum numbers describing each of the two electrons in the 3s orbital. n l m l m s / /2

89 Quantum Numbers: A Macroscale Analogy n l - indicates which train (shell) - indicates which car (subshell) m l - indicates which row (orbital) m s - indicates which seat (spin) No two people can have exactly the same ticket (sit in the same seat).

90 Electron energy states electrons have discrete energy states they fill up the lowest possible energy states in the electron shells and subshells, When all the electrons occupy the lowest possible energies in accord with the foregoing restrictions, an atom is said to be in its ground state. Energy states for a Na atom

91 Electron configurations Most elements: Electron configuration not stable! Element Hydrogen Helium Lithium Beryllium Boron Carbon... Neon Sodium Magnesium Aluminum... Argon... Krypton Atomic # Electron configuration 1s 1 1s 2 (stable) 1s 2 2s 1 1s 2 2s 2 1s 2 2s 2 2p 1 1s 2 2s 2 2p s 2 2s 2 2p 6 (stable) 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 2 3p s 2 2s 2 2p 6 3s 2 3p 6 (stable)... 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 (stable)

92 Electron Configurations Valence electrons those in unfilled shells Filled shells more stable Valence electrons are most available for bonding and tend to control the chemical, electrical, thermal, optical properties example: C (atomic number = 6) 1s 2 2s 2 2p 2 valence electrons

93 Electron Configurations Q. the full electronic configuration of an element is 1s 2 2s 2 2p 5. How many electrons does it have in its outer shell? A. # of outer shell-valence electrons: 7 Q: the full electronic configuration of an element. 1s 2 2s 2 2p 5. What is its atomic number? A. Atomic number: 9

94 Energy Electronic Configurations Fe-atomic # = 26 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2 4d 4p 3d 4s 3p 3s N-shell n = 4 6 electrons left to be located total # e - s 20 electrons M-shell n = 3; 18 electrons 2p 2s 1s L-shell n = 2; 10 electrons K-shell n = 1; 2 electrons

95 Valence electrons Element symbol Atomic number e- configuration H 1 1s 1 1 He 2 1s 2 2 Li 3 1s 2 2s 1 1 Be 4 1s 2 2s 2 2 B 5 1s 2 2s 2 2p 1 3 C 6 1s 2 2s 2 2p 2 4 N 7 1s 2 2s 2 2p 3 5 O 8 1s 2 2s 2 2p 4 6 F 9 1s 2 2s 2 2p 5 7 Ne 10 1s 2 2s 2 2p 6 8 # of valence electrons

96 Order of Subshell Filling The electron configurations of the first ten elements illustrate this point.

97 Electron configurations for common elements

98 Shells and subshells In multi-electron atoms, the energies of the atomic o orbitals are split. Splitting of energy levels refers to the splitting of a shell (n=3) into subshells of different energies (3s, 3p, 3d)

99 Splitting of Shells into subshells 3s subshell 3 rd shell (n = 3p 3; 3) subshell l = 0) (n 3d = subshell 3; l = 1) (n = 3; l = 2) 2s 2 nd subshell shell (n = 2) 2p subshell (n = 2; l = 1) (n = 2; l = 0)

100 Electron Configurations rules for electron configurations: Electrons will reside in the lowest possible energy orbitals Each orbital can accommodate a maximum of two electrons. Electrons will not pair in degenerate orbitals if an empty orbital is available. Orbitals will fill in the order..3p 6 /4s 2 /3d 10 /4p 6 /5s 2 /4d 10 / 5p 6 /6s 2 /4f 14 /5d 10 /6p 6 /7s 2

101 Energy Level Diagram of a multi-electron atom 6s 6p 5d 4f 32 5s 5p 4d 18 4s 4p 3d Arbitrary Energy Scale 3s 3p s 2p 8 1s 2 NUCLEUS

102 Energy Electron Configurations The electron configuration describes how the electrons are distributed in the various atomic orbitals. In a ground state hydrogen atom, the electron is found in the 1s orbital. Ground state electron configuration of hydrogen principal (n = 1) 1s 1 number of electrons in the orbital or subshell 2s 2p 2p 2p angular momentum (l = 0) 1s The use of an up arrow indicates an electron with m s = + ½

103 Energy Electron Configurations If hydrogen s electron is found in a higher energy orbital, the atom is in an excited state. 2s 2p 2p 2p 1s A possible excited state electron configuration of hydrogen 2s 1

104 Energy Electron Configurations Pauli exclusion principle: no two electrons in an atom can have the same four quantum numbers. The ground state electron configuration of helium 2s 2p 2p 2p 1s 2 Quantum number 1s describes the 1s orbital Principal (n) Angular moment (l) Magnetic (m l ) describes the electrons in the 1s orbital Electron spin (m s ) + ½ ½

105 Energy Electron Configurations The Aufbau principle states that electrons are added to the lowest energy orbitals first before moving to higher energy orbitals. Li has a total of 3 electrons The ground state electron configuration of Li 2s 1s 2p 2p 2p The third electron must go in the next available orbital with the lowest possible energy. 1s 2 2s 1 The 1s orbital can only accommodate 2 electrons (Pauli exclusion principle)

106 Energy Electron Configurations Be has a total of 4 electrons 2p 2p 2p 2s 1s The ground state electron configuration of Be 1s 2 2s 2

107 Energy Electron Configurations B has a total of 5 electrons 2p 2p 2p 2s 1s The ground state electron configuration of B 1s 2 2s 2 2p 1

108 Energy Electron Configurations Hund s rule, the most stable arrangement of electrons is the one in which the number of electrons with the same spin is maximized. C has a total of 6 electrons The ground state electron configuration of C 2p 2p 2p 1s 2 2s 2 2p 2 2s 1s The 2p orbitals are of equal energy, or degenerate. Put 1 electron in each before pairing (Hund s rule).

109 Energy Electron Configurations N has a total of 7 electrons 2p 2p 2p 2s 1s The 2p orbitals are of equal energy, or degenerate. Put 1 electron in each before pairing (Hund s rule). The ground state electron configuration of N 1s 2 2s 2 2p 3

110 Energy Electron Configurations O has a total of 8 electrons 2p 2p 2p 2s 1s Once all the 2p orbitals are singly occupied, additional electrons will have to pair with those already in the orbitals. The ground state electron configuration of O 1s 2 2s 2 2p 4

111 Energy Electron Configurations F has a total of 9 electrons 2p 2p 2p 2s 1s When there are one or more unpaired electrons, as in the case of oxygen and fluorine, the atom is called paramagnetic. The ground state electron configuration of F 1s 2 2s 2 2p 5

112 Energy Electron Configurations Ne has a total of 10 electrons 2p 2p 2p 2s 1s When all of the electrons in an atom are paired, as in neon, it is called diamagnetic. The ground state electron configuration of Ne 1s 2 2s 2 2p 6

113 learning check Write the electron configuration and give the orbital diagram of a calcium (Ca) atom (Z = 20). Z = 20, Ca has 20 electrons. Each s subshell can contain a maximum of two electrons, whereas each p subshell can contain a maximum of six electrons. Solution Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Remember that the 4s orbital fills before the 3d orbitals.

114 learning check electron configuration for an arsenic atom (Z = 33) in the ground state. Z = 18 for Ar. The order of filling beyond the noble gas core is 4s, 3d, and 4p. Fifteen electrons go into these subshells because there are = 15 electrons in As beyond its noble gas core Solution As [Ar]4s 2 3d 10 4p Arsenic is a p-block element; therefore, we should expect its outermost electrons to reside in a p subshell.

115 electron configuration? Number of Energy Levels: 3 First Energy Level: 2 Second Energy Level: 8 1s 2 2s 2 2p 6 3s 2 3p 1 Third Energy Level: 3

116 electron configuration? Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 8 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Fourth Energy Level: 1

117 electron configuration? Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 Third Energy Level: 10 Fourth Energy Level: 2

118 electron configuration? Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 13 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Fourth Energy Level: 1

119 electron configuration? Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 13 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 Fourth Energy Level: 2

120 electron configuration? Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 Third Energy Level: 14 Fourth Energy Level: 2

121 electron configuration? Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 18 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 Fourth Energy Level: 1

122 electron configuration? Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 18 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 Fourth Energy Level: 2

123 Valence electrons They occupy the outermost shell. They participate in the bonding between atoms They dictate the physical and chemical properties if the outermost or valence electron shell are completely filled: stable electron configurations occupation of the s and p states for the outermost shell by a total of eight electrons, in neon (Ne), argon (Ar), and krypton (Kr); inert, or noble, gases, which are virtually unreactive chemically.

124 Valence electrons unfilled valence shells assume stable electron configurations by gaining or losing electrons to form charged ions, or by sharing electrons with other atoms. This is the basis for some chemical reactions, and also for atomic bonding in solids.

125 Valence electrons Under special circumstances, the s and p orbitals combine to form hybrid spn orbitals, where n indicates the number of p orbitals involved, which may have a value of 1, 2, or 3. The IIIA, IVA, and VA group elements of the periodic table often form these hybrids. The driving force for the formation of hybrid orbitals is a lower energy state for the valence electrons. For carbon the sp 3 hybrid is of primary importance in organic and polymer chemistries.

126 s- and p-orbitals Aufbau Principle: filling orbitals 1s 2s 2p H: 1s 1 n = 1 l = 0 m l = 0 n = 2 l = 0 m l = 0 m l = -1 n = 2 l = 0 m l = 0 m l = 1

127 s- and p-orbitals Aufbau Principle: filling orbitals 1s 2s 2p He: 1s 2 n = 1 l = 0 m l = 0 n = 2 l = 0 m l = 0 m l = -1 n = 2 l = 0 m l = 0 m l = 1

128 s- and p-orbitals Aufbau Principle: filling orbitals 1s 2s 2p Li: 1s 2 2s 1 n = 1 l = 0 m l = 0 n = 2 l = 0 m l = 0 m l = -1 n = 2 l = 0 m l = 0 m l = 1

129 s- and p-orbitals Aufbau Principle: filling orbitals 1s 2s 2p Be: 1s 2 2s 2 n = 1 l = 0 m l = 0 n = 2 l = 0 m l = 0 m l = -1 n = 2 l = 0 m l = 0 m l = 1

130 s- and p-orbitals Aufbau Principle: filling orbitals B: 1s 2 2s 2 2p 1 1s 2s 2p core closed shell open shell: valence electrons

131 s- and p-orbitals Aufbau Principle: filling orbitals C: 1s 2 2s 2 2p 2 1s 2s 2p Hund s rule: maximum number of unpaired electrons is the lowest energy arrangement.

132 Hund s rule electrons fill orbitals one at a time. we must fill each shell with one electron each before starting to pair them up. the charge of an electron is negative and electrons repel each other. An electron will try to create distance between itself and other electrons by staying unpaired. This further explains why the spins of electrons in an orbital are opposite (i.e. +1/2 and -1/2).

133 s- and p-orbitals Aufbau Principle: filling orbitals N: 1s 2 2s 2 2p 3 O: 1s 2 2s 2 2p 4 1s 2s 2p

134 s- and p-orbitals Aufbau Principle: filling orbitals F: 1s 2 2s 2 2p 5 Ne: 1s 2 2s 2 2p 6 1s 2s 2p

135 s- and p-orbitals Aufbau Principle: filling orbitals Na: 1s 2 2s 2 2p 6 3s 1 or [Ne]3s 1 Mg: 1s 2 2s 2 2p 6 3s 2 or [Ne]3s 2 P: [Ne]3s 2 3p 3 Ar: [Ne]3s 2 3p 6

136 electron configuration? Which one of the following is a proper orbital configuration?

137 electron configuration? Which one of the following is a proper orbital configuration?

138 electron configuration? Which one of the following is a proper orbital configuration?

139 period beyond the d-orbitals s -groups group p -groups 1s2 d-transition elements 2s2/2p6 3s2/3p 6 / 4s 2 /3d 10 /4p 6 5s 2 /4d 10 /5p 6 / 6s 2 /4f 14 /5d 10 /6p 6 lanthanides actinides f-transition elements

140 Organisation of the periodic table

141 Organisation of the periodic table

142 Organisation of the periodic table

143 Organisation of the periodic table

144 Organisation of the periodic table

145 electron configuration? Give electron configurations for the Fe 3+ and S 2- ions. The Fe 3+ ion is an iron atom that has lost three electrons. Since the electron configuration of the Fe atom is 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2, the configuration for Fe3+ is 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5. The S 2- ion a sulfur atom that has gained two electrons. Since the electron configuration of the S atom is 1s 2 2s 2 2p 6 3s 2 3p 4, the configuration for S2- is 1s 2 2s 2 2p 6 3s 2 3p 6.

146 electron configuration? Give the e- configurations for the following ions? Fe 2+ [Ar] 3d 6 4s 2-2 e-: [Ar] 3d 6 Al 3+ [Ne] 3s 2 3p 1-3 e-: [Ne] Cu + [Ar] 3d 10 4s 1-1 e-: [Ar] 3d 10 Ba 2+ [Xe] 6s 2-2 e-: [Xe] Br - [Ar] 3d 10 4s 2 4p 5 +1 e-: [Ar] 3d 10 4s 2 4p 6 O 2- [He] 2s 2 2p 4 +2 e-: [He] 2s 2 2p 6

147 electron configuration? Which of the following electron configurations is an inert gas, a halogen, an alkali metal, an alkaline earth metal, a transition metal? a) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 7 4s 2 b) 1s 2 2s 2 2p 6 3s 2 3p 6 c) 1s 2 2s 2 2p 5 d) 1s 2 2s 2 2p 6 3s 2 e) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 2 4s 2 f) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 transition metal inert gas halogen alkaline earth metal transition metal alkali metal

148 electron configuration? The halogens (group 7A, or group 17, of the periodic table) all have similar chemical properties (for example, forming singly charged negative ions). What aspect of their electron configurations leads to these elements having such similarities? a) They all have a complete 1s 2 shell at the lowest energy level b) They all have an identical Ns 2 Np 5 configuration for their valence electrons (N is any whole number). c) They all have p electrons in their outermost shell. d) They all have an odd number of protons. e) They all have an even number of neutrons.