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Chapter 8 Chapter 8 Evolving model of the atom Chapter 8 ¾ 1803 (Dalton): All matter is composed of tiny, indivisible, indestructible particles called atom. Chapter 8 Atomic Electron Configurations and Chemical Periodicity ¾ 1903 (Thompson): Subatomic particles: electrons and positive charges. Plum-pudding model. ¾ 1911(Rutherford): Protons (positively charge) and neutrons (neutral) are located in the centre of the atom atom. Electrons are somewhere outside the nucleus. ¾ 1913 (Bohr): Electrons are moving in a circular orbit around the nucleus. Only certain orbits with fixed energy are permissible. ¾ORBIT: The circular path in which electrons move around the nucleus ¾ORBITAL: The region in space where an electron is most likely to be found ¾ 1932 (Schrodinger): The region of space (ORBITAL) outside the nucleus where the probability (likelihood) of finding an electron with a given energy is maximum. 1 Chapter 8 ¾ First three quantum numbers (n, l, and ml) describe orbitals shell 1 shell 2 shell 3 shell 4 1s 2s 3s 4s 2p 3p 4p p 3d 4d 4f ¾ shell: Each shell with a designated n has many subshells ¾ subshell: Each subshell with a designated l has many orbitals ¾ orbital: Each orbital with a designated by ml has a specific orientation and has room for TWO electrons Chapter 8 Orbital Energies ¾ What general principle explains orbital energies? ¾ Which orbital has higher energy, 1s, 2s or 3s? Why? Orbital Energies Radial probability Orbitals- Home of Electrons E1s< E2s < E3s Distance from nucleus ¾ Which orbital has higher energy, 2s or 2p? Why? ¾ Which orbital has higher energy, 2px, 2py or 2pz? Why? Radial probability Chapter 8 E2s< E2p What determines the relative energies of these orbitals? Which are lower in energy, which are higher in energy? Distance from nucleus 2 Chapter 8 Chapter 8 Orbital Energies Chapter 8 Effective Nuclear Charge ¾ Zeff: the positive charge actually felt by a valence electron Zeff = Z – s ¾ Z = atomic number ¾ s = shielding parameter ¾ Zeff increases across the p period of p periodic table Effective Nuclear Charge Orbital stability ¾ Lithium ¾ Zeff = 3 – 1.72 = 1.28 ¾ Nitrogen ¾ Zeff = 7 – 3.15 = 3.85 Which electron will be easy to remove, the one from Lithium or Nitrogen? 3 Chapter 8 Chapter 8 Effective Nuclear Charge Orbital stability Effective Nuclear Charge Chapter 8 Magnetic Properties: Electron Orbital stability ¾Zeff: the positive charge actually felt by a valence electron Zeff = Z – s A quantity that comes due to electron-electron repulsion ¾ A physical phenomenon: spinning, charged particles produce magnetic fields ¾ Spinning electrons produce tiny magnetic fields ¾ Electrons can spin in one of two directions 4 Chapter 8 Chapter 8 Magnetic Properties of Electron Chapter 8 The 4th Quantum Number Paired electrons are more stable ¾ Diamagnetic: substances repelled by a strong magnetic field ¾ Paired electrons ¾ Electron spin, ms: ms = ½ or -½ Quantum Mechanical Model and Periodic Table Li ground state Aligned or opposed to the magnetic field ¾ Pauli exclusion principle: ¾ Paramagnetic: substances attracted to a strong magnetic field ¾ Unpaired electrons No two electrons in an atom can have the same set of four quantum numbers n, l, ml, and ms. ¾In order to put more than one electron in an orbital, electrons must have different values of ms. i.e. they must have different spins. ¾Maximum of 2 electrons per orbital 5 Chapter 8 Chapter 8 Energy of Orbitals ¾ For the same type of orbital (same ______), energy increases as n increases (1s < 2s < 3s < 4s…) ¾ For the same n, energy increases s < p < d < f (3s < 3p < 3d) ¾ All orbitals of the same subshell have the same energy (degenerate) (3px = 3py = 3pz) 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p Chapter 8 Energy of Orbitals: n+l rule Orbital Diagrams orbital Draw this diagram and by hand and start filling out electrons. This diagram will be counted as one problem i.e. 1/4th extra credit n=3 3s 3p 3s 3p p 3d subshell n=3 3d shell n=3 3s 3p 3d 6 Chapter 8 Chapter 8 Electron Configuration Rules ¾ Electrons fill the lowest energy orbital first (Aufbau principle) This diagram and any 10 elements’ electron-filled orbital diagram will be counted as one problem i.e. 1/4th extra credit 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f Diagonal Diagram: a guide used to determine the relative energies of subshells in multi-electron atoms Chapter 8 Electron Configuration Rules Energy of Orbitals: Summary ¾Pauli exclusion principle No two electrons in an atom can have the same set of four quantum numbers n, l, ml, and ms. ¾ Two electrons (max) per orbital ¾ Maximize parallel spins when filling a subshell ¾ If more than one orbital in a subshell is available, electrons will fill empty orbitals in the subshell first. (Hund’s Rule) Alternately…. ¾ Electrons preferred to be unpaired as long as an empty orbital with the same energy is available 7 Chapter 8 Chapter 8 Electron Configurations ¾ Three notations for the arrangement of electrons in atoms ¾ Hydrogen Orbital Box Notation ¾ noble gas notation number of electrons Electron Configurations ¾ Hydrogen 1s1 Orbital box diagrams spdf notation Chapter 8 Electron Configurations Lithium # of es =3 Α. 1s22s1 B. 1s12s12p1 C. 2p3 D. 1s3 number of electrons 1s1 orbital type (l) electron shell (n) spdf Notation Orbital Box Notation ¾ Oxygen: # of es =8 Α. 2s22p6 B. 1s12s12p6 C. 1s22s22p4 D. 1s22s32p3 orbital type (l) electron shell (n) spdf Notation 8 Chapter 8 Chapter 8 Chapter 1A 8 Electron Configurations ¾ Hydrogen number of electrons 1s1 Orbital Box Notation ¾ orbital type (l) More Examples ¾ Provide the electron configurations (in orbital box, spdf and noble gas notation) (a) P 8A Transition Metals 2A 3A 4A 5A 6A 7A He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K 3B 4B 5B 6B 7B 8B Ca Sc Ti V Cr Mn Fe 8B 8B 1B 2B Co Ni Cu Zn Ga Ge As Se Br Kr electron shell (n) Rb spdf Notation (b) V Chlorine: # of es =17 Α. 1s22s22p63s23p33d3 B. 1s22s22p63s23p5 C. 1s22s22p53s23p6 D. 1s22s32p63s13p6 H (c) I Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Cs Ba La Hf Sr Y Zr Ta W Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Re Os Ir Pt Au Hg Tl Pb Bi Xe Mt Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr 9 Chapter 8 Chapter 8 Chapter 8 More Examples: Ions Some Anomalies? Periodic Table Organization (a) S2– Chromium and copper So does S2– = Ar? Isoelectronic species ((b)) Br – Transition metal ions (c) Al3+ ¾Half-filled and fully filled d-subshells have extra stability (lower energy). 10 Chapter 8 Chapter 8 Periodic Table Organization atoms where an s subshell is being filled atoms where a p subshell is being filled atoms where a d subshell is being filled ¾ Valence electrons ¾ Core electrons: electrons included in the noble gas notation Li (3): 1s2 2s1 Na(11): 1s22s22p63s1 [He] 2s1 [Ne] 3s1 ¾ Same group = same number and type of valence electrons ¾ Take the case of Li 1s22s1 Radial probability s-block p-block d-block Chapter 8 Effective Nuclear Charge Electron Configurations ¾ Valence electrons: electrons in the outermost shells Î responsible for all macroscopic properties ¾ Core electrons: electrons included in the noble gas notation t ti Li (3): 1s2 2s1 Na(11): 1s22s22p63s1 Distance from nucleus [He] 2s1 [Ne] 3s1 ¾ Same group = same number and type of valence electrons Î Similarity of properties 11 Chapter 8 Chapter 8 Electron Configurations: Atoms and Ions Chapter 8 Periodic Properties ¾ You will need to know the following: Noble gas elements He (2) : 1s2 1. Definitions and chemical equations where appropriate Effective Nuclear Charge ¾ Valence electrons don’t “feel” the full charge of the nucleus ¾ Valence electrons are shielded Ne (10) : [He] 2s2 2p6 Ar (18) : [Ne] 3s2 3p6 Kr (36): [Ar] 4s2 4p6 K+ (19-1= 18) ≡ [Ar] or [Ne] 3s2 3p6 2. Periodic trends moving up and down and left to right across the periodic table ¾ But … valence electrons “feel” a charge that is greater than Z – core electrons ¾ Valence electrons are not completely shielded 3. Explanations of the trends 4. How the atomic properties affect chemical properties Br- (35 +1= 36) ≡ [Kr] or [Ar] 4s2 4p6 12 Chapter 8 Chapter 8 Atomic Size Chapter 8 Atomic Size ¾ The distance from the nucleus to the edge of the outermost electron ¾ Periodic trend: Atomic Size Decrease across a Decrease period Decrease across a across a period period Decrease ¾ Explanation: Effective nuclear charge increases across the group 13 Chapter 8 Atomic Size The best way to explain the increase of atomic size as one goes downward through groups Α. The electrons in a shell repel more, therefore the atom expands B. The nucleus becomes bigger in size as it has more protons and neutrons C Down the group C. group, new shells (i (i.e. e n is increased by 1) are added; each new shell is further and further away from the nucleus D. The nucleus expands and the shells (filled with electrons) expands Chapter 8 Atomic Size The best way to explain the decrease of atomic size as one goes across periods Α. The electrons repel less, therefore the atom shrinks B. The electrons are put on a same shell . The nuclear effective charge increases and the effective pull of the nucleus on its outermost shell electrons increases manyy fold C. Across a period, the total positive charge at the nucleus remains constant D. The nucleus shrinks as it accommodates more neutrons Chapter 8 #1: ¾ Identify the one which is correctly arranged in order of increasing (smallest to largest) atomic size: a. Be, C, O b. Be, O, C c. O, C, Be d. C,O, Be #2: ¾ Identify the one which is correctly arranged in order of increasing (smallest to largest) atomic size: a. Cl, K, S b. Cl, S, K c. K, S, Cl d. K, Cl, S 14 Chapter 8 Ionization Energy (IE) ¾ The energy required to remove an electron from a gaseous atom A(g) + energy Æ A+(g) + e- Chapter 8 Sign Conventions Ionization Energies ¾ Energy absorbed (in) = a positive value + 165 kJ ¾ Energy required (input, raw material) ¾ Energy released (out) = a negative value - 165 kJ ¾ Energy produced (output, product) ¾ Energy input required Decrease Chapter 8 Effective nuclear charge increases across the group The sign tells us which way energy is going The magnitude tells us how much energy is required IE (Be) > IE (B) Be(4): 1s2 2s2 B(5) : 1s2 2s2 2p1 IE (N) > IE (O) N (7): 1s2 2s2 2p3 O (8) : 1s2 2s2 2p4 15 Chapter 8 First Ionization Energy Chapter 8 Chapter 8 Successive Ionizations IE1 IE2 IE3 IE4 IE5 IE6 Na 495 4560 Mg 735 1445 7730 Al 580 1815 2740 Si 780 1575 3220 4350 16100 P 1060 1890 2905 4950 6270 21200 S 1005 2260 3375 4565 6950 8490 Successive Ionizations IE7 ¾ For Mg, 2nd IE > 1st IE ¾ For Al, 3rd IE > 2nd IE > 1st IE ¾ Why? 11600 27000 ¾ For Mg, 3rd IE >>> 2nd IE ¾ For Al, 4th IE >>> 3rd IE Example: Na(g) + IE1 Æ Na+(g) + eNa+(g) + IE2 Æ Na2+(g) + ¾ Why? e- 16 Chapter 8 Chapter 8 Ionization Energies: Summary ¾First ionization energies generally increase across a period and decrease down a group ¾Effective nuclear charge increases across the th group #3: ¾ Arrange each set of atoms in increasing IE1: a. Sr, Ca, Ba b. Ba, Sr, Ca c. Ca, Sr, Ba d. Ba, Ca, Sr #4: ¾ Arrange each set of atoms in increasing IE1: a. Br, Rb, Se b. Br, Se, Rb c. Rb, Br, Se d. Rb, Se, Br Chapter 8 Electron Affinity ¾ The energy released when an electron is added to a gaseous atom A(g) + e- Æ A-(g) + energy ¾ A free electron is not a stable. It would always be associated with an atom. 17 Chapter 8 Chapter 8 Electron Affinity Predictions Chapter 8 Electron Affinity Trends Electron Affinity Summary A(g) + e- Æ A-(g) + energy Exception ¾ Across a period: Should it get easier or harder to add an electron? ¾ Down a group: Should it get easier or harder to add an electron? ¾ If it’s easy to add an electron, is the EA a large negative number or a small negative number? ¾ Deviations from the general trends ¾An element with a high ionization energies generally has a high affinity for an electron. ¾Effective nuclear charge increases across the group and decreases down a group 18 Chapter 8 Chapter 8 Trends in Metallic Behavior Relative tendencies to lose and gain electrons Acid-base Behaviors of Elemental OXides Chapter 8 Ionization: Change in Size ¾ Why does the size decrease? Metals donate electrons to oxygen Nonmetals share electrons to oxygen 3 p+ and 3 e3 p+ and 2 e- Ionic Metal oxides react with water to produce hydroxides (OH-) that are basic Covalent Nonmetal oxides react with water to produce acids that releases proton in solution H+ Elements at the left form cations easily Elements at the right form anions easily 19 Chapter 8 Chapter 8 Ionization: Change in Size ¾ Why does the size increase? 9 p+ and 10 e9 p+ and 9 e- Review ¾ Zeff: the positive charge actually felt by a valence electron ¾ Atomic size: The distance from the nucleus to the edge of the outermost electron ¾ IE: The energy required to remove an electron from a gaseous atom. ¾ Successive ionization ¾ EA: The energy released when an electron is added to a gaseous atom ¾ Ion sizes Chapter 8 The Reaction of Na and Cl Na IE 495 4560 EA EA > 0 Cl 1251 -348 ¾ How can we use these numbers to explain the product of the reaction? ¾ Is NaCl2 a reasonable product? ¾ Is Na2Cl a reasonable product? 20 Chapter 8 Chapter 8 Periodic trends and Chemical Properties • Reactivity of metals Chemical Reactivity Summary ¾ Noble gases high IE, low EA do not react ¾ Metals low IE, low EA lose electrons ¾ Non-metals high IE, high EA add electrons • Reactivity of nonmetals ¾ Metal + non-metal Æ metal loses e-’s and non-metal gains e-’s ¾ non-metal + non-metal Æ shared e-’s 21