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Edit Comment Close Premium member Presentation Transcript Slide 1: Chapter 03 Electronic Configuration And Chemical Periodicity Slide 2: Mendeleev said: “ ... the elements arranged according to their atomic weights, show a distinct periodicity (regular variation) of their properties” 3.1 The Periodic Law and the Periodic Table 1869 Dmitri Mendeleev Lothar Meyer Slide 3: The Periodic Table helps us to Predict the electronic structure and physical and chemical properties of the variuos elements Understand different aspects of chemical bonding 3.1 The Periodic Law and the Periodic Table Slide 4: 3.1 The Periodic Law and the Periodic Table Slide 5: Group or Family (Vertical) Period (Horizontal) 3.1 The Periodic Law and the Periodic Table The Periodic table : The Periodic table Slide 7: Metal - is a substance whose atoms tend to lose electrons during chemical change, forming positive ions. Nonmetal – is a substance whose atoms may gain electrons, forming negative ions. 3.1 The Periodic Law and the Periodic Table Slide 8: Metals Good conductors of heat and electricity. Malleable and ductile. Moderate to high melting points. Nonmetals Nonconductors of heat and electricity. Brittle solids. Some are gases at room temperature. 3.1 The Periodic Law and the Periodic Table Slide 9: Atomic Number and Atomic Mass 3.1 The Periodic Law and the Periodic Table Slide 10: Electronic Configuration – describes the arrangement of electrons in atoms. Valence Electrons – are the outermost electrons in an atom, which are involved, or have the potential to become involved in the bonding process. 3.2 Electron Arrangement and the Periodic Table Slide 11: 3.2 Electron Arrangement and the Periodic Table Slide 12: 3.2 Electron Arrangement and the Periodic Table Slide 13: 3.2 Electron Arrangement and the Periodic Table Example 3.1 Provide the total number of electrons, total number of valence electrons, and energy level in which the valence electrons are found for the silicon (Si) atom. Total number of electrons = 14 Total number of valence electrons = 4 Energy level (n) = 3 Slide 14: Quantum Numbers Principal Energy Level (n) Describes the size and energy level of the orbital Commonly called shell Positive integer (n = 1, 2, 3, 4, …) As the value of n increases: The energy increases The average distance of the e- from the nucleus increases Angular-Momentum Quantum Number (l) Commonly called subshell There are n different shapes for orbitals If n = 1 then l = 0 If n = 2 then l = 0 or 1 If n = 3 then l = 0, 1, or 2 etc. Commonly referred to by letter (subshell notation) l = 0 s (sharp) l = 1 p (principal) l = 2 d (diffuse) l = 3 f (fundamental) 3.2 Electron Arrangement and the Periodic Table Slide 15: Quantum Numbers Magnetic Quantum Number (ml ) Defines the spatial orientation of the orbital There are 2l + 1 values of ml and they can have any integral value from -l to +l If l = 0 then ml = 0 If l = 1 then ml = -1, 0, or 1 If l = 2 then ml = -2, -1, 0, 1, or 2 etc. 3.2 Electron Arrangement and the Periodic Table Slide 16: 3.2 Electron Arrangement and the Periodic Table Slide 17: 3.2 Electron Arrangement and the Periodic Table Slide 18: - Electrons have spin which gives rise to a tiny magnetic field and to a spin quantum number (ms). 3.2 Electron Arrangement and the Periodic Table Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers. Slide 19: 3.2 Electron Arrangement and the Periodic Table Hund’s Rule: If two or more orbitals with the same energy are available, one electron goes into each until all are half-full. The electrons in the half-filled orbitals all have the same spin. Aufbau Principle: Lower-energy orbitals fill before higher-energy orbitals. Slide 20: Electronic Configuration Obtain the total number of electrons in the atoms from the atomic number found on the periodic table Fill up the orbitals with electrons following Pauli’s Exclusion Principle, Hund’s Rule of Multiplicity, and Aufbau Building Up Principle 3.2 Electron Arrangement and the Periodic Table Example 3.2 Give the electronic configuration of Tin (Sn) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p2 Slide 21: Shorthand Electron Configuration Na = 1s22s22p63s1 Ne = 1s22s22p6 Na = [Ne] 3s1 Work out the shorthand electron configuration for Mg-12, Cl-17, and K-19. Mg = [Ne] 3s2 Cl = [Ne] 3s23p5 K = [Ar] 4s1 3.2 Electron Arrangement and the Periodic Table Slide 22: 3.3 The Octet Rule Octet Rule – Elements usually react in such a way as to obtain the electron configuration of the noble gas closest to them in the periodic table. Ion Formation and the Octet Rule Metals tend to form cations Na Na+ + e- Mg Mg2+ + 2e- Al Al3+ + 3e- Na+, Mg2+, and Al3+ are all isoelectronic Isoelectronic – elements which have the same number of electrons Nonmetal tend to form anions F + e- F- O + 2e- O2- N + 3e- N3- Transition metals tend to have variable valence Fe has Fe2+ and Fe3+ Cu has Cu+ and Cu2+ Slide 23: 3.4 Trends in the Periodic Table Atomic Radius Atomic Radius – size of an atom Slide 24: 3.4 Trends in the Periodic Table Electron Affinity Electron Affinity(Eea) - is a measure of the energy change when an electron is added to a neutral atom to form a negative ion. X(g) + e- X-(g) Slide 25: Size of Ions 3.4 Trends in the Periodic Table Slide 26: 3.4 Trends in the Periodic Table Ionization Energy Ionization Energy (I or EI) - is the energy required to remove one mole of electrons from one mole of isolated gaseous atoms or ions X X+ + e- Slide 27: 3.4 Trends in the Periodic Table Electronegativity Electronegativity (En) - is a chemical property that describes the ability of an atom (or, more rarely, a functional group) to attract electrons (or electron density) towards itself in a covalent bond HF HBr HCl HI Slide 29: 3.4 Trends in the Periodic Table Trends in the Periodic Table You do not have the permission to view this presentation. In order to view it, please contact the author of the presentation.