Transcript Chapter 8
Chapter 7 The Structure of Atoms and Periodic Trends Arrangement of Electrons in Atoms Electrons in atoms are arranged as: Shells (n) Subshells (l) Subshell orientation (ml) Pauli’s Exclusion Principle discovered in 1925 by Wolfgang Pauli -No two electrons in an atom can have the same set of 4 quantum numbers Practice: What are the 4 quantum numbers for each electron in He? Aufbau Principle Describes the electron filling order in atoms -electrons are placed in the lowest available energy orbital -the periodic table is a function of electron configurations for the elements Electron Configuration To remember the correct filling order for electrons in atoms: Electron Configuration Writing Electron Configurations Two ways to express electron configuration: 1. spdf notation Example: H atomic number = 1 1s 1 no. of electrons value of l value of n Writing Electron Configurations 2. Orbital box notation ORBITAL BOX NOTATION for He, atomic number = 2 spdf notation 2 1s 1s Arrows depict electron spin Electron Configurations Using the Aufbau Principle to determine the electronic configurations of the elements 1st row elements: 1s 1 H 2 He Configuration 1 1s 2 1s Electron Configurations Hund’s rule: electrons fill suborbitals by placing electrons in each suborbital unpaired first with the same spin direction, then the electrons pair Electron Configurations 3s 3p Configuration 11 Na Ne 12 Mg Ne 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar Ne Ne Ne Ne Ne Ne Ne 3s1 Ne 3s 2 Ne 3s 2 3p1 Electron Configurations and Quantum Numbers We can write a complete set of quantum numbers for all of the electrons in every element: – Na – Ca – Fe 3s 11 Na Ne 3p Configuration Ne 3s1 Electron Configurations and Quantum Numbers n l ml 1 0 0 2 nd e - 1 0 0 3rd e - 2 0 0 4 th e - 2 0 0 5th e- 2 1 6 th e - 2 1 7 th e - 2 1 8th e- 2 1 9 th e - 2 1 10th e - 2 1 11th e - 3 0 1st e - ms 1/2 1 s elect rons 1/2 1/2 2 s elect rons 1/2 1/2 0 1/2 1 1/2 2 p elect rons 1 1/2 The ml and ms are interchangeable 0 1/2 1 1/2 0 1/23 s elect ron -1 Electron Configurations and Quantum Numbers Noble Gas Notation (or short hand notation): The first 18 electrons in Ca are represented with the preceding noble gas ([Ar]) - we only concern ourselves with the outermost e- 3d 20 Ca [Ar] 4s 4p Configuration Ar 4s2 Skip the first 18 electrons Electron Configurations and Quantum Numbers n l ml [Ar]19th e - 4 0 0 20th e - 4 0 0 ms 1/2 4 s elect rons 1/2 Electron Configurations and Quantum Numbers There is only one set of 4 quantum numbers for each of the 26 electrons in Fe: – To save space, we use the symbol [Ar] to represent the first 18 electrons in Fe 3d 26 Fe Ar 4s 4p Configuration Ar 4s2 3d6 Electron Configurations of Ions Electrons are removed from subshell of highest energy level (n-level) P0 [Ne] 3s2 3p3 -3e- ---> P3+ [Ne] 3s2 3p0 3p 3p 3s 3s 2p 2p 2s 2s 1s 1s Electron Configurations of Ions For transition metals, remove the highest s-orbital electrons first: Fe [Ar] 4s2 3d6 -2 electrons Fe2+ [Ar] 3d6 -3 electrons Fe3+ [Ar] 3d5 To form cations, always remove electrons of highest n value first! More About the Periodic Table Representative Elements Groups IA, IIA, IIIA-VIIIA – These elements will have their “outermost” electron in an outer s or p orbital – Variations in their properties are similar from top-tobottom More About the Periodic Table d-Transition Elements All have d electrons -With n s-orbitals -With n-1 d–orbitals Have small property variations from row-torow More About the Periodic Table f - transition metals -Sometimes called inner transition metals -Electrons are being added to f orbitals Extremely small variations in properties from one element to another More About the Periodic Table Noble Gases -Have filled electron shells -have similar chemical reactivities -similar electronic structures He Ne Ar Kr Xe Rn 1s2 [He] 2s2 2p6 [Ne] 3s2 3p6 [Ar] 4s2 4p6 [Kr] 5s2 5p6 [Xe] 6s2 6p6 Periodic Properties • Atomic radii describes the relative sizes of atoms • Atomic radii increase within a column • Atomic radii decrease within a row Periodic Properties Example: Arrange these elements based on their atomic radii: Se, S, O, Te O < S < Se < Te Periodic Properties Example: Arrange these elements based on their atomic radii: P, Cl, S, Si Cl < S < P < Si Periodic Properties Electronegativity: measure of the tendency of an atom to attract electrons to itself -Fluorine is the most electronegative element -Cesium is the least electronegative element Electronegativity increase from left-to-right and decrease from top-to-bottom increase decrease Periodic Properties Example: Arrange these elements based on their electronegativity: Se, Ge, Br, As Ge < As < Se < Br Periodic Properties Example: Arrange these elements based on their electronegativity: Be, Mg, Ca, Ba Ba < Ca < Mg < Be Periodic Properties Ionization Energy: energy required to remove an electron from an atom in the gas state First ionization energy (IE1) – Energy required to remove the first electron from an atom in the gas state to form a 1+ ion Atom(g) + energy Atom+(g) + eExample: Mg(g) + 738kJ/mol Mg+ + e- Periodic Properties Second ionization energy (IE2) – The amount of energy required to remove the second electron from a gaseous 1+ ion Atom+ + energy Atom2+ + eMg+ + 1451 kJ/mol Mg2+ + e- Atoms can have 3rd (IE3), 4th (IE4), etc. - Each IE is significantly higher than the previous IE Periodic Properties Ionization Energy: 1. IE2 > IE1 always takes more energy to remove a second electron from an ion 2. IE1 increases to the right Important exceptions are Be & Mg, N & P, etc. due to filled and half-filled subshells 3. IE1 decrease down First Ionization Energies He 2500 Ne 2000 Ionization Energy (kJ/mol) N 1500 1000 H C Be F Ar Cl P O Mg S B 500 Li Ca Si Na Al K 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 Atomic Number Periodic Properties Example: Arrange these elements based on their first ionization energies: Sr, Be, Ca, Mg Sr < Ca < Mg < Be Periodic Properties Example: Arrange these elements based on their first ionization energies: Al, Cl, Na, P Na < Al < P < Cl Periodic Properties Electron Affinity: Energy absorbed when an electron is added to an atom to form a negative ion Sign conventions for electron affinity: – If electron affinity > 0 energy is absorbed – If electron affinity < 0 energy is released Electron affinity is the measure of an atom’s ability to form negative ions atom(g) + e- + EA atom-(g) Periodic Properties Examples of electron affinity values: Mg(g) + e- + 231 kJ/mol Mg-(g) EA = +231 kJ/mol Br(g) + e- Br-(g) + 323 kJ/mol EA = -323 kJ/mol Increasing ability to decreasing ability to add electrons add electrons Electron Affinity Electron Affinities of Some Elements Electron Affinity (kJ/mol) 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 0 -50 -100 -150 -200 -250 -300 -350 -400 He Be B N Ne Mg Al Ar P Na H Li K O C Si S F Atomic Number Ca Cl Periodic Properties Example: Arrange these elements based on their electron affinities: Al, Mg, Si, Na Si < Al < Na < Mg Periodic Properties Ionic Radius: diameter of an atom in its ionized form -Cations are always smaller Element Atomic Radius (Å) Ion Ionic Radius (Å) Li Be 1.52 1.12 Li+ Be2+ 0.90 0.59 Periodic Properties Anions are always larger Element N O F Atomic Radius(Å) Ion 0.75 0.73 0.72 N3- O2- F1- Ionic Radius(Å) 1.71 1.26 1.19 Periodic Properties Cation radii decrease from left to right across a period – Increasing nuclear charge attracts the electrons and decreases the radius. Ion Rb+ Sr2+ In3+ Ionic Radii(Å) 1.66 1.32 0.94 Periodic Properties Anion radii decrease from left to right across a period – Increasing electron numbers in highly charged ions cause the electrons to repel and increase the ionic radius Ion N3- O2- F1- Ionic Radii(Å) 1.71 1.26 1.19 Ionic Radii Active Figure 8.15 Periodic Properties Example: Arrange these elements based on their ionic radii: Ca2+, K+, Ga3+ K1+ > Ca2+ > Ga3+ Periodic Properties Example: Arrange these elements based on their ionic radii: Cl-1, Se-2, Br-1, S-2 Cl1- < S2- < Br1- < Se2-