118 known elements All elements 1 (hydrogen) to 118 (ununoctium) have been discovered or synthesized. Of these, all up to and including californium (98) exist naturally; the rest have only been synthesized in laboratories. Chapter 4, Figure 4.1, On the periodic table, groups are the elements arranged as vertical columns, and periods are the elements in each horizontal row.

Au, (aurum) Fe, iron (ferrum) Ag, (argentum) Elements and Symbols A symbol Represents the name of an element (118 of them). Consists of 1 or 2 letters. Starts with a capital letter. 1-Letter Symbols 2-Letter Symbols C carbon Co cobalt N nitrogen Ca calcium F fluorine Al aluminum O oxygen Mg magnesium Some symbols are derived from Latin names are: Au, (aurum) Fe, iron (ferrum) Ag, (argentum) 2

Some groups are known by common names. Periodic Table 7 periods 18 groups Some groups are known by common names. 3 3

Nonmetals are located to the right. Are dull, brittle, and poor conductors of heat and electricity. Are good insulators. Metalloids - located along the heavy zigzag line between the metals and nonmetals. Are better conductors than nonmetals, but not as good as metals. Are used as semiconductors and insulators. Figure 03-06 Title: Metalloids Caption: Along the heavy zigzag line on the periodic table that separates the metals and nonmetals are metalloids, which exhibit characteristics of both metals and nonmetals. Metals are located to the left. Are shiny and ductile. Are good conductors of heat and electricity. 4

elements  atoms  subatomic particles The Atom elements  atoms  subatomic particles location electrical charge Protons at the nucleus have a positive (+) charge. Neutrons at the nucleus are neutral. Electrons at a distance from the nucleus have a negative (-) charge.

The number of protons is fixed. Atomic Number and Mass Number The number of protons is fixed. the number of protons = The atomic number Is specific for each element. In the periodic table it appears above the symbol of an element. Chapter 3, Unnumbered Figure, Page 95

Isotopes The number of neutrons can vary 1 1 𝐻 1 2 𝐻 1 3 𝐻 1 𝐻 2 𝐻 3 𝐻 protons + neutrons = The mass number Since ps and ns are the heavier particles (e- is very light), they are the main contributors to the mass of an atom. The mass number is used in atomic symbols, where all nuclear particles (ps and ns) are stated: Isotopes Atoms of the same element that vary in the number of neutrons (mass numbers). 1 1 𝐻 1 2 𝐻 1 3 𝐻 1 𝐻 2 𝐻 3 𝐻 H-1 H-2 H-3

Isotopes and Atomic Mass The average mass of all isotopes of an element based on their % abundance in nature, “weighted average”. 8

Practice 1. Match the elements to the description: a. Metals in Group 4A(14) 1) Sn, Pb 2) C, Si 3) C, Si, Ge, Sn b. Nonmetals in Group 5A(15) 1) As, Sb, Bi 2) N, P 3) N, P, As, Sb c. Metalloids in Group 4A(14) 1) C, Si, Ge, 2) Si, Ge 3) Si, Ge, Sn, Pb 2. An atom of zinc has a mass number of 65. a. How many protons are in this zinc atom? 1) 30 2) 35 3) 65 b. How many neutrons are in the zinc atom? 1) 30 2) 35 3) 65 c. What is the mass number of a zinc atom that has 37 neutrons? 1) 37 2) 65 3) 67 3. An atom has 14 protons and 20 neutrons. a. Its atomic number is 1) 14 2) 16 3) 34 b. Its mass number is 1) 14 2) 16 3) 34 c. The element is 1) Si 2) Ca 3) Se 4. Write the nuclear symbols for atoms with the following subatomic particles: a. 8 p+, 8n, 8e- ___________ b. 17p+, 20n, 17e- ___________ 9

Electron ENERGY LEVELS The energy of the electron has a minimum value and all other energy values of it are integer multiples (n) of this minimum. E = n (min. value) n= positive integer (1,2,3,4…) Analogy: Money has a minimum value of 1 penny ($0.01) and all other coins or bills are restricted to integer values of this minimum: $ = n (0.01) n = a positive integer (1,2,3,4…) Example: you can hand someone 75 ($0.01)=$0.75 but not 75.5 ($0.01)=$0.755 Max Planck The name given to this minimum value is quantum and we say that energy is quantized. Therefore, the energy of an e- can only have a specific value based on the positive integer, n. n=4 Energy increases as n increases Lowest energy at n = 1 (closest to nucleus) n=3 n=2 max number of e-s = 2n2 n=1 2 n=2 8 n=3 18 n=4 32 n=1 Neils Bohr nucleus

E ∝ 1  Electromagnetic energy = radiant energy Propagates as a wave Thermal energy = heat Electromagnetic energy = radiant energy Propagates as a wave wavelength E ∝ 1  If the wavelength of this radiation lies in the visible range, we see a colored light.

Electrons in lowest available energy level = ground state. By absorbing E, the e- is raised to a higher energy level called the excited state. An e- loses energy when it falls to a lower energy level and emits a photon (E carrier). The photon travels in space as a wave. The energy of the wave is inversely proportional to its wavelength. E∝ 1 λ

364.6 nm 656.3 nm 656.3 nm 656.3 nm

f d d p p p s s s s ENERGY SUBLEVELS In order to accommodate all the e-s, energy levels are subdivided into energy sublevels. These are identified by the letters s, p, d, and f. The number of sublevels = n n=1 1 sublevel n=2 2 sublevels n=3 3 sublevels f d d nucleus p p p s s s s n=1 n=2 n=3 n=4

Orbitals An orbital Is a three-dimensional space around a nucleus where an electron is most likely to be found. Generally denoted by a box. The number of orbitals = n2 n=1 12 =1 n=2 22 =4 n=3 32 =9 7 5 3 1 Erwin Schrödinger Physicist f d d p p p nucleus s s s s n=1 n=2 n=3 n=4

Can hold up to 2 electrons. These must spin in opposite directions. Orbitals Can hold up to 2 electrons. These must spin in opposite directions. Arrows are used to represent these e-s.

Practice 5. In each of the following energy level changes, indicate if energy is: 1) absorbed 2) emitted 3) not changed. a. An electron moves from the first energy level (n = 1) to the third energy level (n = 3). b. An electron falls from the third energy level to the second energy level. c. An electron moves within the third energy level. 6. Indicate the number and type of orbitals in each of the following: a. 4s sublevel b. 3d sublevel c. n = 3 7. The number of A. electrons that can occupy a p orbital is 1) 1 2) 2 3) 3 B. p orbitals in the 2p sublevel is 1) 1 2) 2 3) 3 C. d orbitals in the n = 4 energy level is 1) 1 2) 3 3) 5 D. electrons that can occupy the 4f sublevel is 1) 2 2) 6 3) 14 8. Indicate the type and number of orbitals in each of the following energy levels or sublevels: a.) 3p sublevel b. ) n = 2 c.) n = 3 d.) 4d sublevel 17 LecturePLUS Timberlake 17

Electron Configurations An electron configurations is a shorthand notation describing an electron position in an atom. Take C as an example: total of 6 e- 14 10 6 2 f 2 d d nucleus 2 2 p p p s s s s n=1 n=2 n=3 n=4 Energy level 18

4s lower than 3d due to sublevel overlap The Aufbau (“Build-up”) Principle (from the German aufbauen) The order of filling various electron subshells with electrons follows the same order given by the arrows in this diagram. Iron 26 e- 1s2 2s2 2p6 3s2 3p6 4s2 3d6 energy level sublevel number of electrons 4s lower than 3d due to sublevel overlap 1s 2s 2p 3s 3p 3d 4s Ar 1s2 2s2 2p6 3s2 3p6 K 1s2 2s2 2p6 3s2 3p6 4s1 Ca 1s2 2s2 2p6 3s2 3p6 4s2 Sc 1s2 2s2 2p6 3s2 3p6 3d1 4s2 Ti 1s2 2s2 2p6 3s2 3p6 3d2 4s2 19

Exceptions to the Aufbau order In describing the ground state electron configuration the guiding idea is that it corresponds to an isolated atom in its lowest total energy. 1. The orbital energies depend on a number of factors such as nuclear charge and interactions of electrons in different occupied orbitals. 2. The energy scale varies with atomic number. In other words, all atoms of a given element have the same set of energy levels, but atoms of different elements have different sets. 3. Some orbitals are very close together, so their order can change, depending on the occupancies of other orbitals. The only exceptions you are asked to remember are those based on the special stability of filled or half-filled d-orbitals.

Abbreviated Configurations An abbreviated configuration shows The symbol of the noble gas in brackets that represents completed sublevels. The remaining electrons in order of their sublevels, Example: Chlorine has a configuration of 1s2 2s2 2p6 3s2 3p5 The abbreviated configuration for chlorine is [Ne] 3s2 3p5 Iron 1s2 2s2 2p6 3s2 3p6 4s2 3d6 [Ar] 4s2 3d6 21 21

C 1s2 2s2 2p2 Orbital Diagrams diagram Orbitals as boxes. diagram Electrons as vertical arrows. C 1s2 2s2 2p2 The Pauli Exclusion Principle Electrons in the same orbital with opposite spins (up and down vertical arrows). Hund’s rule Fill orbitals in sublevels of the same type with one electron until half full, then pair up using opposite spins. Wolfgang Pauli Friedrich Hund

Hund’s rule

Sublevel Blocks 1s He O Ca O 1s2 2s2 2p4 Ca 1s2 2s2 2p6 3s2 3p6 4s2

Periodic Trends The valence electrons are the electrons in the outermost, highest energy level are related to the Group number of the element determine the chemical properties of the elements Example: Phosphorus has 5 valence electrons. P Group 5A(15) 1s2 2s22p6 3s23p3 n =1 n=2 n=3

An electron-dot symbol Convenient way of representing valence electrons using dots around the symbol of the element. · · · Mg· or Mg· or · Mg or · Mg · The same for all members in a group. · Be · · Mg · · Ca · · Sr · · Ba · Groups 1A(1) to 4A(14) use single dots · · Na · · Mg · · Al · · C· · Groups 5A(15) to 7A(17) use pairs and single dots. · · · · · P· : O· · ·

Practice 9. Write the electron configuration and abbreviated configuration for the following elements: a. S b. K 10. Write the orbital diagrams for a. nitrogen b. oxygen 11. Use the sublevel blocks on the periodic table to write the electron configuration for bromine. 27

Practice  14. a. X is the electron-dot symbol for 1) Na 2) K 3) Al 12. State the number of valence electrons for each. a. Calcium b. Group 6A (16) c. Tin 13. State the number of valence electrons for each. a. 1s22s22p63s23p1 b. 1s22s22p63s2 c. 1s22s22p5  14. a. X is the electron-dot symbol for 1) Na 2) K 3) Al   b.  X  is the electron-dot symbol for 1) B 2) N 3) P 15. Write the electron-dot symbol for each of the following elements: bromine b. aluminum