Atomic Structure - Nanyang Girls' High School Chemistry PDF

Nanyang Girls’ High School Name: Science Department (Chemistry) Class: Atomic Structure Learning Outcomes: Students should be able to: U M (a) identify and describe protons, neutr...

Nanyang Girls’ High School Name: Science Department (Chemistry) Class: Atomic Structure Learning Outcomes: Students should be able to: U M (a) identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses (b) describe, with the aid of diagrams, the structure of an atom as consisting of protons and neutrons (nucleons) in the nucleus and electrons arranged in shells (energy levels) the distribution of mass and charges within an atom define proton (atomic) number and nucleon (mass) number interpret and use nuclide notations such as (c) deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (d) describe the contribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number (e) define isotopes and distinguish between isotopes based on the different numbers of neutrons present (f) describe the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals (to be covered in Term 3) (g) describe the shapes of s, p and d orbitals (to be covered in Term 3) (h) state the electronic configuration of atoms given the proton number (to be covered in Term 3) U: I have understood the LO M: I have mastered the LO 1 Bohr model of atom electron shell nucleus comprising protons (p) and neutrons (n) electron (e) Electrons orbiting around nucleus at certain energy levels (electron shells) Structure of atom Subatomic particle Symbol Relative charge Relative mass Proton p +1 1 Neutron n 0 1 electron e –1 Atoms are electrically neutral (no. of protons = no. of electrons) A An atom can be represented using nuclide notation Z X nucleon (mass) number Total no. of nucleons (protons + neutrons) in the nucleus A Z X, where X is the symbol of the element proton (atomic) number Total no. of protons Isotopes Definition: Isotopes are atoms of the same element with the same number of protons, but different number of neutrons. Same chemical properties (due to same number of electrons) Slightly different physical properties (due to different number of neutrons) 2 Practice 1 particle number of protons number of neutrons number of electrons A 8 10 8 B 11 7 10 C 9 10 10 D 8 8 8 Which 2 particles are isotopes? A, D Which particle(s) are ions? B, C Write the nuclide notation for particle A or Relative atomic mass, Ar Definition: The average mass of one atom of an element compared to the mass of an atom of carbon-12. Ar = Calculation of Ar Example: isotope % abundance chlorine-35 75 chlorine-37 25 Calculate the relative atomic mass of chlorine. Ar of chlorine = 35 x 0.75 + 37 x 0.25 = 35.5 [Note: Ar is to one decimal place] Practice 2a: A newly discovered element X contains 3 major isotopes in the following abundances: 5.9% of 54 X, 91.8% of 56 X and 2.3% of 57 X. Calculate the relative atomic mass of X. Relative atomic mass = 0.059 x 54 + 0.918 x 56 + 0.023 x 57 = 55.9 3 Calculation of Relative Abundance Example: Boron has 2 isotopes with relative masses of 10 and 11. Given that boron has a relative atomic mass of 10.8, calculate the relative abundance of boron–10 and boron–11. Let the relative abundance of boron-10 be x 10x + 11(1 – x) = 10.8 10x + 11 – 11x = 10.8 x = 0.2 Therefore relative abundance of boron–10 is 20.0% and relative abundance of boron–11 is 80.0%. [Note: calculated final values are generally to 3 significant figures] Practice 2b: Given that gallium contains 2 isotopes with relative masses of 69 and 71, calculate the relative abundance of gallium–69 and gallium–71. Let the relative abundance of gallium-69 be x 69x + 71(1 – x) = 69.7 (Ar from Periodic Table) x = 0.65 Therefore relative abundance of gallium–69 is 65.0% and relative abundance of gallium–71 is 35.0%. 4  Quantum Model (Electron Cloud Model) of an Atom nucleus electron cloud (orbital) Developed by many scientists, including Heisenberg, de Broglie, Schrodinger Electrons occupy regions of space known as orbitals. o An atomic orbital is a region of space with > 95% probability of finding an electron. 5 Electronic structure of atoms (covered in Term 3) Principal quantum number (n) Principal quantum number describes the energy level and size (radius) of an atomic orbital The larger n is, the higher the energy level and the further away it is from the nucleus. This is similar to the electron shell model learnt previously. Subshells Each principal quantum number has a fixed number of subshells i.e the type of orbitals present at that energy level The number of subshells for any given principal quantum number is equal to the numerical value of the principal quantum number. Energy levels of subshells within a principal quantum number increase in the following order: s < p < d < f Atomic orbitals Each type of subshell has one or more orbitals having the same energy but different orientations in space. Each orbital contains 2 electrons with opposite spins. Principal Number of quantum orbitals for Maximum number of electrons in each Subshells shell/ respective quantum shell (2n2) number, n subshells 1 1s 1 2 2 2s, 2p 1, 3 8 3 3s, 3p, 3d 1, 3, 5 18 4 4s, 4p, 4d, 4f 1, 3, 5, 7 32 6 Shapes of orbitals (covered in Term 3) s orbital spherical shape size increases as principal quantum number increase 1s 2s 3s p orbital dumbbell shape size increases as principal quantum number increases p orbitals in the same quantum shell have the same size and shape but different orientation d and f orbitals (Enrichment) 7  Rules for writing electronic configuration (covered in Term 3) Aufbau (“building up”) Principle Electrons occupy orbitals in order of energy levels, ie orbital with lowest energy is filled first. Pauli Exclusion Principle Each orbital can have at most 2 electrons. Electrons in an orbital must have opposite spins. ↿ Hund’s rule When filling subshells with more than 1 orbital of the same energy (p, d, f), orbitals must be singly occupied before electrons are paired. example: nitrogen - 1s2 2s2 2p3 no electron-electron repulsion electron-electron repulsion  lower energy (more stable)  higher energy (less stable) ↿ ↿ ↿ ↿ ↿ ↿ ↿ ↿ ↿ 1s2 2s2 2p3 1s2 2s2 2p3   8  Atom No. of spdf notation Electron Orbital Diagram electrons 1 1s1 1 H 1 1s 2s 2p 3s 3p 3d 4s 4p 4d 4 1s2 2 He 2 1s 2s 2p 3s 3p 3d 4s 4p 4d 7 1s 2s 2 1 3 Li 3 1s 2s 2p 3s 3p 3d 4s 4p 4d 9 1s2 2s2 4 Be 4 1s 2s 2p 3s 3p 3d 4s 4p 4d 11 1s2 2s2 2p1 5 B 5 1s 2s 2p 3s 3p 3d 4s 4p 4d 12 1s2 2s2 2p2 6C 6 1s 2s 2p 3s 3p 3d 4s 4p 4d 14 1s 2s 2p 2 2 3 7N 7 1s 2s 2p 3s 3p 3d 4s 4p 4d 16 1s2 2s2 2p4 8 O 8 3p 4s 4p 4d 1s 2s 2p 3s 3d 19 1s 2s 2p 2 2 5 9F 9 1s 2s 2p 3s 3p 3d 4s 4p 4d 9  Atom No. of spdf notation Electron Orbital Diagram electrons 20 1s2 2s2 2p6 10 Ne 10 1s 2s 2p 3s 3p 3d 4s 4p 4d 39 1s2 2s2 2p6 3s2 3p6 4s1 19 K 19 1s 2s 2p 3s 3p 3d 4s 4p 4d 40 1s2 2s2 2p6 3s2 3p6 4s2 20 Ca 20 1s 2s 2p 3s 3p 3d 4s 4p 4d 45 1s 2s 2p 3s 3p 3d 4s 2 2 6 2 6 1 2 21 Sc 21 1s 2s 2p 3s 3p 3d 4s 4p 4d 48 1s2 2s2 2p6 3s2 3p6 3d2 4s2 22Ti 22 1s 2s 2p 3s 3p 3d 4s 4p 4d 52 1s2 2s2 2p6 3s2 3p6 3d5 4s1 24 Cr 24 1s 2s 2p 3s 3p 3d 4s 4p 4d 55 1s2 2s2 2p6 3s2 3p6 3d5 4s2 25 Mn 25 1s 2s 2p 3s 3p 3d 4s 4p 4d 64 1s2 2s2 2p6 3s2 3p6 3d10 4s1 29 Cu 29 1s 2s 2p 3s 3p 3d 4s 4p 4d 65 1s 2s 2p 3s 3p 3d 4s 2 2 6 2 6 10 2 30 Zn 30 1s 2s 2p 3s 3p 3d 4s 4p 4d 10  Relate to arrangement of Periodicity and electronic configurations elements in Periodic Table From the electronic configurations you wrote earlier, what did you notice about the valence electrons for the various groups in the periodic table? Fill in the table below. Group Valence electronic configuration 1 ns1 2 ns2 13 ns2np1 14 ns2np2 15 ns2np3 16 ns2np4 17 ns2np5 18 ns2np6 Practice a) Write the electronic configuration, in spdf notation, for P. a) P: 1s2 2s2 2p6 3s2 3p3 b) Write the electronic configuration, in spdf notation, for Co, and fill in the following energy level diagram. b) Co: 1s2 2s2 2p6 3s2 3p6 3d7 4s2 11  Annex (Enrichment) Enrichment: Electronic configurations of ions Write electronic configurations for atom first For cations: remove number of electrons equal to charge for cations, starting with the outermost orbital (not necessarily the orbital with the highest energy level) e.g. Electrons are removed from the 4s orbitals first, rather than the 3d orbitals. For anions: add number of electrons equal to charge for anions, using Aufbau Principle, Pauli Exclusion Principle, Hund’s Rule (see page 8) Examples Ion Electronic configuration Ion Electronic configuration Na+ 1s2 2s2 2p6 Cu2+ 1s2 2s2 2p6 3s2 3p6 3d9 O2– 1s2 2s2 2p6 F– 1s2 2s2 2p6 Fe2+ 1s2 2s2 2p6 3s2 3p6 3d6 Al3+ 1s2 2s2 2p6 Enrichment: Reasons for anomalous cases (Cr and Cu) 3d and 4s are very close in energy levels Inter-electronic repulsion is minimised by placing electron in 3d instead of 4s (for Cr) Change in the energy of orbitals across the period (for Cu) o Across the period from left to right, 3d orbitals may have a lower energy level than 4s. Hence, 3d orbitals will be filled first. Z = atomic number 12