Evolving model of the atom

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1 Atomic Electron Configurations and Chemical Periodicity Evolving model of the atom 1803 (Dalton): All matter is composed of tiny, indivisible, indestructible particles called atom (Thompson): Subatomic particles: electrons and positive charges. Plum-pudding model. 1911(Rutherford): Protons (positively charge) and neutrons (neutral) are located in the centre of the atom. Electrons are somewhere outside the nucleus (Bohr): Electrons are moving in a circular orbit around the nucleus. Only certain orbits with fixed energy are permissible (Schrodinger): The region of space (ORBITAL) outside the nucleus where the probability (likelihood) of finding an electron with a given energy is maximum. ORBIT: The circular path in which electrons move around the nucleus ORBITAL: The region in space where an electron is most likely to be found 1

2 Orbitals- Home of Electrons Orbital Energies Orbital Energies First three quantum numbers (n, l, and m l ) describe orbitals shell 1 1s shell 2 2s 2p shell 3 3s 3p 3d shell 4 4s 4p 4d 4f shell: Each shell with a designated n has many subshells subshell: Each subshell with a designated l has many orbitals orbital: Each orbital with a designated by m l has a specific orientation and has room for TWO electrons What general principle explains orbital energies? Which orbital has higher energy, 1s, 2s or 3s? Why? Which orbital has higher energy, 2s or 2p? Why? Which orbital has higher energy, 2p x, 2p y or 2p z? Why? Radial probability Radial probability Distance from nucleus E 1s < E 2s < E 3s E 2s < E 2p What determines the relative energies of these orbitals? Which are lower in energy, which are higher in energy? Distance from nucleus 2

3 Orbital Energies Effective Nuclear Charge Z eff : the positive charge actually felt by a valence electron Z eff = Z s Z = atomic number s = shielding parameter Z eff increases across the period of periodic table eff p p Effective Nuclear Charge Orbital stability Lithium Z eff = = 1.28 Nitrogen Z eff = = 3.85 Which electron will be easy to remove, the one from Lithium or Nitrogen? 3

4 Effective Nuclear Charge Orbital stability Effective Nuclear Charge Orbital stability Z eff : the positive charge actually felt by a valence electron Z eff = Z s A quantity that comes due to electron-electron repulsion Magnetic Properties: Electron A physical phenomenon: spinning, charged particles produce magnetic fields Spinning electrons produce tiny magnetic fields Electrons can spin in one of two directions 4

5 Magnetic Properties of Electron Paired electrons are more stable Diamagnetic: substances repelled by a strong magnetic field Paired electrons The 4 th Quantum Number Electron spin, m s : m s = ½ or -½ Aligned or opposed to the magnetic field Quantum Mechanical Model and Periodic Table Li ground state Paramagnetic: substances attracted to a strong magnetic field Unpaired electrons Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers n, l, m l, and m s. In order to put more than one electron in an orbital, electrons must have different values of m s. i.e. they must have different spins. Maximum of 2 electrons per orbital 5

6 Energy of Orbitals Energy of Orbitals: n+l rule Orbital Diagrams For the same type of orbital (same ), energy increases as n increases (1s < 2s < 3s < 4s ) For the same n, energy increases s < p < d < f (3s < 3p < 3d) All orbitals of the same subshell have the same energy (degenerate) (3p x = 3p y = 3p z ) 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p 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 orbital 3s 3p 3d n=3 subshell 3s 3p 3d n=3 shell 3s 3p 3d n=3 6

7 Electron Configuration Rules Electron Configuration Rules Energy of Orbitals: Summary Electrons fill the lowest energy orbital first (Aufbau principle) This diagram and any 10 elements electron-filled orbital 1s diagram will be counted as one problem i.e. 1/4th extra credit 2s 2p 3s 3p 3d Diagonal Diagram: 4s 4p 4d 4f a guide used to 5s 5p 5d 5f determine the 6s 6p 6d 6f relative energies 7s 7p 7d 7f of subshells in multi-electron atoms Pauli exclusion principle No two electrons in an atom can have the same set of four quantum numbers n, l, m l, and m s. Two electrons (max) per orbital Maximize parallel spins when filling a subshell If more than one orbital in a subshell is available, electrons will fill empty orbitals in the subshell first. (Hund s Rule) Alternately. Electrons preferred to be unpaired as long as an empty orbital with the same energy is available 7

8 Electron Configurations Electron Configurations Electron Configurations Three notations for the arrangement of electrons in atoms Hydrogen 1s 1 number of electrons Hydrogen 1s 1 number of electrons Orbital box diagrams Orbital Box Notation orbital type (l) Orbital Box Notation orbital type (l) spdf notation noble gas notation Lithium #ofes=3 Α. 1s 2 2s 1 B. 1s 1 2s 1 2p 1 C. 2p 3 D. 1s 3 electron shell (n) spdf Notation Oxygen: #ofes=8 Α. 2s 2 2p 6 B. 1s 1 2s 1 2p 6 C. 1s 2 2s 2 2p 4 D. 1s 2 2s 3 2p 3 electron shell (n) spdf Notation 8

9 Electron Configurations Hydrogen number of electrons 1s 1 Orbital Box Notation orbital type (l) electron shell (n) spdf Notation More Examples Provide the electron configurations (in orbital box, spdf and noble gas notation) (a) P (b) V Chapter 1A 8 H 2A Li Be Transition Metals Na Mg 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A 5A 6A 7A He B C N O F Ne Al Si K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn P S Cl I 8A Ar Xe Chlorine: # of es =17 Α. 1s 2 2s 2 2p 6 3s 2 3p 3 3d 3 B. 1s 2 2s 2 2p 6 3s 2 3p 5 C. 1s 2 2s 2 2p 5 3s 2 3p 6 D. 1s 2 2s 3 2p 6 3s 1 3p 6 (c) I Fr Ra Ac Rf Db Sg Bh Hs Mt Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np PuAmCmBk Cf Es Fm Md No Lr 9

10 Some Anomalies? More Examples: Ions Periodic Table Organization Chromium and copper (a) S 2 So does S 2 = Ar? (b) Br Isoelectronic species Transition metal ions (c) Al 3+ Half-filled and fully filled d-subshells have extra stability (lower energy). 10

11 s-block p-block d-block Periodic Table Organization Valence electrons atoms where an s subshell is being filled atoms where a p subshell is being filled atoms where a d subshell is being filled Core electrons: electrons included in the noble gas notation Li (3): 1s 2 2s 1 Na(11): 1s 2 2s 2 2p 6 3s 1 [He] 2s 1 [Ne] 3s 1 Same group = same number and type of valence electrons l probability Radia Effective Nuclear Charge Distance from nucleus Take the case of Li 1s 2 2s 1 Electron Configurations Valence electrons: electrons in the outermost shells responsible for all macroscopic properties Core electrons: electrons included in the noble gas notation Li (3): 1s 2 2s 1 Na(11): 1s 2 2s 2 2p 6 3s 1 [He] 2s 1 [Ne] 3s 1 Same group = same number and type of valence electrons Similarity of properties 11

12 Electron Configurations: Atoms and Ions Noble gas elements He (2) : 1s 2 Ne (10) : [He] 2s 2 2p 6 Ar (18) : [Ne] 3s 2 3p 6 Kr (36): [Ar] 4s 2 4p 6 K + (19-1= 18) [Ar] or [Ne] 3s 2 3p 6 Br - (35 +1= 36) [Kr] or [Ar] 4s 2 4p 6 Periodic Properties You will need to know the following: 1. Definitions and chemical equations where appropriate 2. Periodic trends moving up and down and left to right across the periodic table 3. Explanations of the trends 4. How the atomic properties affect chemical properties Effective Nuclear Charge Valence electrons don t feel the full charge of the nucleus Valence electrons are shielded But valence electrons feel a charge that is greater than Z core electrons Valence electrons are not completely shielded 12

13 Atomic Size Atomic Size Atomic Size The distance from the nucleus to the edge of the outermost electron Periodic trend: Decrease Decrease across a Decrease period across a period Decrease across a period Explanation: Effective nuclear charge increases across the group 13

14 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, new shells (i.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 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 many fold C. Across a period, the total positive charge at the nucleus remains constant D. The nucleus shrinks as it accommodates more neutrons #1: 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: 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

15 Ionization Energy (IE) Sign Conventions Ionization Energies The energy required to remove an electron from a gaseous atom Energy input required A(g) + energy A + (g) + e - Energy absorbed (in) = a positive value kj Energy required (input, raw material) Energy released (out) = a negative value kj Energy produced (output, product) The sign tells us which way energy is going The magnitude tells us how much energy is required Decrease Effective nuclear charge increases across the group IE (Be) > IE (B) Be(4): 1s 2 2s 2 B(5) : 1s 2 2s 2 2p 1 IE (N) > IE (O) N (7): 1s 2 2s 2 2p 3 O (8) : 1s 2 2s 2 2p 4 15

16 First Ionization Energy Successive Ionizations Successive Ionizations IE 1 IE 2 IE 3 IE 4 IE 5 IE 6 IE 7 Na Mg Al Si P S Example: Na(g) + IE 1 Na + (g) + e - Na + (g) + IE 2 Na 2+ (g) + e - For Mg, 2 nd IE > 1 st IE For Al, 3 rd IE > 2 nd IE > 1 st IE Why? For Mg, 3 rd IE >>> 2 nd IE For Al, 4 th IE >>> 3 rd IE Why? 16

17 Ionization Energies: Summary First ionization energies generally increase across a period and decrease down a group Effective nuclear charge increases across the group #3: 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: Arrange each set of atoms in increasing IE1: a. Br, Rb, Se b. Br, Se, Rb c. Rb, Br, Se d. Rb, Se, Br Electron Affinity The energy released when an electron is added to a gaseous atom A(g) + e - A - (g) + energy A free electron is not a stable. It would always be associated with an atom. 17

18 Electron Affinity Predictions Electron Affinity Trends Electron Affinity Summary A(g) + e - A - (g) + energy Across a period: Should it get easier or harder to add an electron? Exception Down a group: Should it get easier or harder to add an electron? If it s easy to add an electron, is the EA a large negative number or a small negative number? Deviations from the general trends An element with a high ionization energies generally has a high affinity for an electron. Effective nuclear charge increases across the group and decreases down a group 18

19 Trends in Metallic Behavior Acid-base Behaviors of Elemental OXides Ionization: Change in Size Relative tendencies to lose and gain electrons Metals donate electrons to oxygen Nonmetals share electrons to oxygen Why does the size decrease? 3 p + and 3 e - 3 p + and 2 e - Ionic Covalent Metal oxides react with water to produce hydroxides (OH - ) that are basic 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

20 Ionization: Change in Size Why does the size increase? 9 p + and 10 e - 9 p + and 9 e - Review Z eff : the positive charge actually felt by a valence electron Atomic size: The distance from the nucleus to the edge of the outermost electron IE: The energy required to remove an electron from a gaseous atom. Successive ionization EA: The energy released when an electron is added to a gaseous atom Ion sizes The Reaction of Na and Cl IE EA Na 495 EA > Cl How can we use these numbers to explain the product of the reaction? Is NaCl 2 a reasonable product? Is Na 2 Cl a reasonable product? 20

21 Periodic trends and Chemical Properties Reactivity of metals Chemical Reactivity Summary Noble gases high IE, low EA do not react Metals low IE, low EA lose electrons Non-metals high IE, high EA add electrons Reactivity of nonmetals Metal + non-metal metal loses e - s and non-metal gains e - s non-metal + non-metal shared e - s 21

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