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Topic 8 Periodic Properties of the Elements PowerPoint PPT Presentation

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Topic 8 Periodic Properties of the Elements. Electron Configuration. An “electron configuration” of an atom is a particular distribution of electrons among available subshells.

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Topic 8 Periodic Properties of the Elements

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Topic 8

Periodic Properties of the Elements

Electron Configuration

An “electron configuration” of an atom is a particular distribution of electrons among available subshells.

  • The notation for a configuration lists the subshell symbols (s, p, d, f) sequentially with a superscript indicating the number of electrons occupying that subshell.

  • For example, lithium (atomic number 3) has two electrons in the “1s” sub shell and one electron in the “2s” sub shell 1s2 2s1.

Electron Configuration

  • Ground state configuration refers to the most stable (lowest energy) way of assigning the electrons.

  • All other ways are higher energy or excited state configurations that are unstable and caused by absorbing light or collisions with other particles.

  • Excited atoms eventually revert to their ground state configuration either by giving off light or as a result of collisions with other particles.

  • If ground or excited state is not specified, we assume the term electron configuration refers to the ground state.

The Pauli Exclusion Principle

The maximum number of electrons allowed per subshell based on Pauli’s Exclusion Principle:

Aufbau Principle

  • Listed below is the order in which all the possible sub-shells fill with electrons.

  • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f

  • The Aufbau principle is a scheme used to reproduce the ground state electron configurations of atoms by following the “building up” order.

    lowest energy


    • This order is based on increasing energy levels using the following factors:

    • Orbitals with higher (n + l values) have higher energies

    • i.e. 3d (3+2=5) is higher energy than 4s (4+0=4)

      • because the n+l value is higher (note: d, l = 2 and s, l=0).

    • For orbitals with the same (n + l values), the orbital with the higher n have higher energies

    • i.e. 3s (3+0=3) has the same n+l value as 2p (2+1=3) but 3s is higher energy because of the higher n of 3.

    Configurations and the Periodic Table

    • We can use the periodic table as a memory aid for the typical order of filling orbitals.

    p block – 3 orbital with max 2e- each = 6 electrons ; p6 max

    s block – 1 orbital with max 2e- = 2 electrons; s2 max

    d block – 5 orbital with max 2e- each = 10 electrons; d10 max

    f block – 7 orbital with max 2e- each = 14 electrons;

    f14 max

    • To write the ground state electron configuration of an atom, you first need to locate the atom on the periodic table. For a neutral atom, the number of electrons equals the number of protons (atomic number). You will write the subshells with the number of electrons in the subshell written as a superscript (nl#e-) from left to right and top to bottom until you reach the particular atom in question on the periodic table. Basically, each block represents an electron that you are adding to the configuration.

    • i.e. (21 electrons to account for in electron config.)








    Ground state electron config. for Sc



    • i.e. (5 electrons to account for in ground state electron config.)




    • i.e. (10 electrons to account for in ground state electron config.)




    • i.e. (11 electrons to account for in ground state electron config.)





    • i.e. (82 electrons to account for in ground state electron config.)
















    • There is a short notation for writing electron configurations which involves using the noble gas core. We specify the noble gas in the row above the atom in question in brackets representing the electron configuration to that point and then write the configuration of the remaining electrons. For the above Pb example, we would write:







    Configurations and the Periodic Table

    • There are exceptions to the order of filling the electron configuration: notable exceptions are Cr, Cu, Nb, Ru, Rh, Pd, Ag, Gd, and Pt.

    • The commonly used justification for these exceptions is that the electrons are “shifted” in order to have half-filled (d5) or fully-filled (d10)

    • There seems to be some stability attained with half-filled or fully-filled subshells.

    • But this justification is tenuous since the exceptions are not consistent.

    i.e. Cu: [Ar] 3d10 4s1 is the actual electron configuration instead of [Ar] 3d9 4s2 as predicted; one electron has shifted from the “s” orbital to the “d” level for the stability gained from having the inner shell completely filled.

    Valence Electrons

    Valence Electronsare electrons that reside in the outermost shell of an atom - or in other words, those electrons outside the “noble gas core”.

    • These electrons are primarily involved in chemical reactions.

    • Elements within a given group (vertical column of periodic table) have the same “valence shell configuration.”

    • This accounts for the similarity of the chemical properties among groups of elements in the same family (vertical group).

    P: [Ne] 3s23p3

    As: [Ar] 4s2 3d10 4p3note: 3d electrons are inner electrons and not valence electrons.

    Sb: [Kr] 5s2 4d105p3

    These three atoms reside in the same VA group and have the same valence electrons: 2 in the s block and 3 in the p block.

    Valence Electrons

    Valence electrons for the groups:

    Group IA: ns1

    Group IIA: ns2

    Group IIIA: ns2 np1

    Group IVA: ns2 np2

    Group VA: ns2 np3

    Group VIA: ns2 np4

    Group VIIA: ns2 np5

    Group VIIIA: ns2 np6

    Transition metals: ns2 (n-1)dx where x = 1 to 10

    i.e. Sc: [Ar] 4s2 3d1

    Configurations and the Periodic Table

    • So far we have done ground state electron configurations for neutral atoms. What happens if you have an ion (charged atom)?

    • Since we are counting electrons, any charge will either add or subtract the number of electrons we are placing in the electron configuration meaning the number of electrons does not equal the number of protons.

    • For cations, we will subtract electrons from the total.

    • For anions, we will add electrons to the total.

      • Ca2+subtract 2e- because you have lost 2e-

      • Cl-add 1e- because you have gained 1e-

    Configurations and the Periodic Table

    Ca: 1s2 2s2 2p6 3s2 3p64s2

    Here are a few examples:

    Ca2+: 1s2 2s2 2p6 3s2 3p6lost two outer shell electrons

    Cl: 1s2 2s2 2p6 3s2 3p5

    Cl-: 1s2 2s2 2p6 3s2 3p6gained one electron

    Note that Ar (1s2 2s2 2p6 3s23p6), Ca2+, and Cl- all have the exact same electron configurations. We say that these species are isoelectronic with each other. They have the same arrangement of electrons but are not the same species; same number of electrons but different number of protons.

    Electrons do not come off the same way they go on; the outer most valence electrons come off before any inner electrons;

    V: [Ar] 4s2 3d3

    V3+: [Ar] 3d2 lost three electrons; the 2 outermost 4s and 1 inner 3d. Note we did not lose all 3 in d level.

    Orbital diagrams

    An orbital diagram gives more detailed information than an electron configuration. In an orbital diagram,

    • electrons are represented by arrows

    • boxes, circles, or blanks are used to represent orbitals

    • upward-pointing arrows represent electrons with +1/2 spin and downward-pointing arrows represent electrons with -1/2 spin

    • Pauli’s principle must be followed meaning no more than 2 electrons per orbital with opposite spin.

    • We fill the orbitals in the highest occupied subshell singly with electrons of the same spin before putting a second electron (of opposite spin) in any orbital. This is known as Hund’s Rule of Maximum Multiplicity. The justification for Hund’s rule is spin correlation; electrons with same spin repel each other less than electrons with opposite spin.

    • Each group of orbitals is labeled by its subshell notation.

    1s 2s 2p

    Orbital diagrams

    • Each orbital is represented by a box, circle or a line.

    An orbital diagram is used to show how the orbitals of a subshell are occupied by electrons.

    • Electrons are represented by arrows:

    • up for ms= +1/2 and down for ms= -1/2

    • We fill the orbitals in the highest occupied subshell singly with electrons of the same spin before putting a second electron (of opposite spin) in any orbital.

    Neon: 10e-

    Orbital Diagrams

    Here’s another way orbital diagrams are displayed by increasing energy levels. This example is for Vanadium with 23 electrons:









    Orbital Diagrams

    The maximum number of electrons and their orbital diagrams are:

    Paramagnetic and Diamagnetic Substances

    Although an electron behaves like a tiny magnet, two electrons that are opposite in spin cancel each other. Only atoms with unpaired electrons exhibit magnetic properties.

    Paramagnetic substance– is one that is weakly attracted by a magnetic field as the result of one or more unpairedelectrons. Whenever the electron configuration of an atom has partially filled subshells, the atom is paramagnetic. Vanadium on slide 17 had unpaired electrons in the 3d level; therefore, is paramagnetic.

    Diamagnetic substance– is not attracted by a magnetic field because it has only paired electrons. Neon on slide 16 had all electrons paired; therefore, is diamagnetic.

    HW 61-62

    code: electron

    Periodic Properties

    Recall that periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically (hence the name periodic table).

    • We will look at three periodic properties:

      • Atomic radius

      • Ionization energy

      • Electron affinity

    Periodic Properties

    • Covalent radius – of atom is ½ distance between the nuclei of two like atoms joined in a molecule.

    • ionic radius same as above.

    • Within each period (horizontal row), the atomic radius tends to increase with decreasing atomic number (right to left).

    Atomic radius (size)

    • Within each group (vertical column), the atomic radius tends to increase with the period number (top to bottom).


    Two factors determine the size of an atom:

    • Why increase down?

    • As we go down the periodic table, the valence shell gets larger. We are adding more shells as we go down the table making the outer electrons farther from the nucleus with less attraction for the nucleus.

    • Why decrease across?

    • The effective nuclear charge, which is the positive charge an electron experiences from the nucleus minus any “shielding effects” from intervening electrons, gets larger as we go across the periodic table. As the effective nuclear charge increases, valence electrons will be pulled toward the nucleus more and held more tightly making the atom smaller.

    Figure: Representation of atomic radii (covalent radii) of the main-group elements.

    Ex. Arrange the following in order of increasing atomic radius:

    Ca, N, As, F, Ba

    increasing means smallest atomic radius to largest radius



    Locating the atoms on the periodic table and comparing them, we find



    Atomic size periodic trend increases across from right to left and top to bottom on the table


    < As

    < Ca

    < Ba (largest size)

    < N

    (smallest size) F

    Ionic Radii

    Cations are smaller than neutral atom: Na+ < Na

    Anions are larger than neutral atom: Cl- > Cl

    Cation has more protons than electrons; therefore electrons are pulled more towards the nucleus making it smaller

    Cation lost 3s shell when lost e-; therefore, cation smaller

    Anion has more electrons than protons; therefore, electrons are pulling more outwards from nucleus making it larger. There is more electron-electron repulsion and a smaller effective nuclear charge causing the anion to get larger.

    Within an isoelectronic group of ions, the one with the greatest nuclear charge (largest excess of protons) will be the smallest.

    Ex. Ar, Cl-, Ca2+, S2-, K+all have 18e-; same electron configuration but different number of protons

    Z= 18, 17, 20, 16, 19

    For isoelectronic groups, the larger the positive charge (higher effective nuclear charge), the more pull the valence electrons will have by the nucleus, the smaller the species.

    < K+



    < S2- largest

    smallest Ca2+

    Ionization energy = 520 kJ/mol

    Periodic Properties

    • The first ionization energy of an atom is the minimal energy needed to remove the highest energy (outermost) electron from the neutral atom in gaseous state.

    Ionization energy

    • For a lithium atom, the first ionization energy is illustrated by:

    Metal atoms tend to give up valence electrons while nonmetal tend to gain electrons. Energy must be absorbed (endothermic process) to remove an electron; therefore, we have +ionization energies.

    Periodic Properties


    • There is a general trend that ionization energies increase across a period (horizontal row); this trend is the opposite of atomic radius because it is more difficult to remove an electron that is closer to the nucleus in smaller atoms.

    Ionization energy (IE)

    • The trend for ionization energies increases up a group (vertical column).

    • Summary: ionization energy increases towards the top right corner of the periodic table with He having the highest ionization energy (larger the radius, smaller the IE)

    • Note there are some exceptions to the trend such as between groups IIA and IIIA and VA and VIA.

    Periodic Properties

    • The electrons of an atom can be removed successively.

    Ionization energy

    • The energies required at each step are known as the first ionization energy, the second ionization energy, and so forth.

    • 2nd IE tends to be larger than 1st IE because you are dealing with ions (charged species) which make it harder to pull an e- from a charged ion than a neutral species; therefore, more energy will be required.

    2nd Ionization energy = 7298 kJ/mol

    1st Ionization energy = 520 kJ/mol

    Ex. Arrange the following in order of increasing 1st ionization energy:

    As, Ca, O, N

    increasing means smallest IE to largest IE



    Locating the atoms on the periodic table and comparing them, we find



    IE periodic trend increases across and up the table

    (smallest IE) Ca

    < N

    < O(largest IE)

    < As

    Electron Affinity = -349 kJ/mol

    Periodic Properties

    Electron Affinity

    • The electron affinity is the energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion.

    • Electron affinity process releases energy (exothermic process); therefore, tend to be negative values.

    Periodic Properties

    Electron Affinity

    • The more negative the electron affinity, the more stable the negative ion that is formed.

    • Broadly speaking, the general trend goes from lower left to upper right as electron affinities become more negative.

    increases in negative value

    HW 63

    code: trend

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