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Announcements

  • Audio 0:00:39.475037
  • Scores got changed up, but now they’re changed back to what they were originally
  • Ch 4 homework due next week

Electron Configuration and Elemental Properties: The Metals

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  • Metallic elements make up the majority of the elements in the periodic table.
    • Alkali Metals:
      • They have one more electron than the previous noble gas and occupy the first column.
      • In their reactions, the alkali metals lose one electron, and the resulting electron configuration is the same as that of a noble gas.
        • Forming a cation with a 1+ charge
    • Alkaline Earth Metals:
      • They have two more electrons than the previous noble gas and occupy the second column.
  • In their reactions, the alkaline earth metals lose two electrons, and the resulting electron configuration is the same as that of a noble gas.
    • Forming a cation with a 2+ charge

Clicker 1

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  • Give the number of core electrons for Cd
    • Cd atomic # is 48, but based on it’s position on the periodic table, the d, valence electrons are not core, so the answer is 46

Metallic Behavior and Electron Configuration

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Gaining or Losing Electrons

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  • Electrons want to be in the lowest lying orbital that has a vacancy
  • If that is on another atom: electron transfers and get an ion
  • Orbitals shrink across periodic table as nuclear charge increases across periodic table
  • For metals: energy of orbitals is (generally) higher than surrounding matter: lose electrons and make cations
  • For non-metals: energy of orbital is (generally) lower than surrounding matter: gain electrons and make anions

Orbital Blocks and Their Position in the Periodic Table

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Electron Configuration and Elemental Properties: Noble Gases

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  • The noble gases have eight valence electrons.
    • Except for He, which has only two electrons
  • They are especially nonreactive.
    • He and Ne are practically inert.
  • The reason the noble gases are so nonreactive is that the electron configuration of the noble gases is especially stable.

Electron Configuration and Ion Formation

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  • Ion formation can be predicted by an element’s location in the periodic table.
  • These atoms form ions that will result in an electron configuration that is the same as that of the nearest noble gas.
  • Metals form cations (positively charged atoms).
    • Alkali metals (group 1A) form only +1 cations.
    • Alkaline earth metals (group 2A) form only +2 cations.
    • Transition, inner transition, and p-block metals form a variety of charged cations.
  • Nonmetals form anions (negatively charged atoms).
  • Halogens (group 7A) usually gain one electron to form –1 anions.
    • Other nonmetals can form a variety of charged anions.

Electron Configuration and Ion Formation: Elements that Form Ions with Predictable Charges

Effective Nuclear Charge and the Screening Effect

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    • Nuclear charge minus core electrons

Effective Nuclear Charge

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  • The effective nuclear charge is a net positive charge that is attracting a particular electron.
  • Core electrons efficiently shield electrons in the outermost principal energy level from nuclear charge.
  • Outermost electrons in the valence shell do not efficiently shield one another from nuclear charge.
  • Z is the nuclear charge, and S is the number of electrons in lower energy levels.
    • Electrons in the same energy level contribute to screening, but since their contribution is so small, they are not part of the calculation.
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  • There are several methods for measuring the radius of an atom,and they give slightly different numbers.
  • Van der Waals radius = nonbonding
  • Covalent radius = bonding radius
  • Atomic radius is an average radius of an atom based on measuring large numbers of compounds.
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  • Atomic radius decreases across a period (left to right). – Adding electrons to the same valence shell
    • Effective nuclear charge increases.
    • Valence shell held closer
  • Atomic radius increases down a group.
    • Valence shell farther from nucleus
    • Effective nuclear charge fairly constant (expansion mostly due to increasing n)

Summarizing Atomic Radii Trend for Main- Group Elements

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  • The size of an atom is related to the distance the valence electrons are from the nucleus.
    • The larger the orbital an electron is in, the farther its most probable distance will be from the nucleus, and the less attraction it will have for the nucleus.
    • Traversing down a group adds a principal energy level, and the larger the principal energy level an orbital is in, the larger its volume.
    • Quantum-mechanics predicts that the atoms should get larger down a column.

Summarizing Atomic Radii Trend for Main- Group Elements

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  • The larger the effective nuclear charge an electron experiences, the stronger the attraction it will have for the nucleus.
    • The closer their average distance will be to the nucleus, the stronger the attraction the valence electrons have for the nucleus.
    • Traversing across a period increases the effective nuclear charge on the valence electrons.
    • Quantum-mechanics predicts that the atoms should get smaller across a period.

Summarizing Atomic Radii Trend for Transition Elements

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  • Atoms in the same group increase in size down the column.
  • Atomic radii of transition metals are roughly the same size across the d block.
    • Much less difference than across main group elements
    • Valence shell ns2, not the (n − 1)d electrons
    • Effective nuclear charge on the ns2 electrons approximately the same

Ions: Magnetic Properties

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  • Electron configurations that result in unpaired electrons mean that the atom or ion will have a net magnetic field; this is called paramagnetism.
    • Will be attracted to a magnetic field
  • Electron configurations that result in all paired electrons mean that the atom or ion will have no magnetic field; this is called diamagnetism.
    • Slightly repelled by a magnetic field

Radii of Atoms and Their Ions: Cations

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  • Cation radius is smaller than its corresponding atom radius.
    • Loose electrons experiencing small effective nuclear charge; remaining electrons those experiencing a larger effective nuclear charge
  • Traversing down a group increases the (n − 1) level, causing the cations to get larger.
  • Traversing to the right across a period increases the effective nuclear charge for isoelectronic cations, causing the cations to get smaller.
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  • isoelectronic = same number of electrons

Radii of Atoms and Their Ions: Anions

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  • When atoms form anions, electrons are added to the valence shell.
  • Addition of electrons increases repulsion in valence shell without compensating increase in effective nuclear charge
  • The result is that the anion is larger than the atom. Traversing down a group increases the n level, causing the anions to get larger.
  • Traversing to the left across a period decreases the effective nuclear charge for isoelectronic anions, causing the anions to get larger.
  • oposite of what happens with cations

Clicker 2

  • Place the following in order of increasing radius
    • Br-, Na+, Rb+
    • Na+, Rb+, Br-

Periodic Trend: Ionization Energy (Potential)

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  • Ionization Energy (IE):
  • It is the minimum energy needed to remove an electron from an atom or ion in the gas phase.
  • It is an endothermic process (requires the input of energy to remove the electron)
    • Valence electron easiest to remove, lowest IE
  • First ionization energy = energy to remove electron from neutral atom – All atoms have first ionization energy. M(g) + IE1 èM1+(g) + 1 e–
  • Second IE = energy to remove from 1+ ion, etc. M+1(g)
  • IE2 è M2+(g) + 1 e–
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    • increases as you go across the periodic table
    • Charge shrinks the nucleus, takes more energy to remove the charge

Periodic Trend: Ionization Energy (Potential)

  • Audio 0:40:38.726742 Ionization Energy (IE):
  • The larger the effective nuclear charge on the electron to be removed, the more energy it takes to remove it.
  • The farther the most probable distance the electron is from the nucleus, the less energy it takes to remove it.
  • Trend:
    • First IE decreases down the group.
  • Valence electron is farther from nucleus.
    • First IE generally increases across the period. + Effective nuclear charge increases.

Periodic Trend: Ionization Energy (Potential)

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    • Nitrogen has three unpaired electrons, the next electron has to go down an orbital which will make it a little easier to remove the last electron in oxygen to get to nitrogen.

First Ionization Energy: Exceptions to the Trend

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  • GENERAL trend for first ionization energy of main-group elements is that as you go across a period, ionization energy increases.
    • Exceptions: 2A to 3A and 5A to 6A
  • Exceptions are usually a result of
    • the type of orbital (s, p, d, or f) and its shielding ability;
    • repulsion factors associated with electrons occupying degenerate orbitals (i.e., p orbitals).
  • B has smaller first ionization energy than Be due to electron position: 2p for B and 2s for Be. The electron in 2p orbitals has more shielding (i.e., lower effective nuclear charge) and therefore requires less energy for its removal than an electron in a 2s orbital.
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  • They depend on the number of valence electrons an element has. – ionization energies increase dramatically in going from valence to core electrons
  • Removal of each successive electron costs more energy. – Shrinkage in size due to having more protons than electrons – Outer electrons closer to the nucleus; therefore harder to remove
  • There’s a regular increase in energy for each successive valence electron.

Clicker 3

What period 3 element has the following ionization energies (all in kJ/mol)? IE1 = 1012 IE2 = 1900 IE3= 2910 IE4= 4960 IE5= 6270 A) Si B) S C) P D) Cl E) Mg IE6 = 22,200

Vocab

Term Definition
alkali metals elements which have one more electron than the previous noble gas and occupy the first
alkaline Earth Metals elements which have two more electrons than the previous noble gas and occupy the second column
metalloids an element whose properties are between those of metals and solid nonmetals (the “stairs” on the table)
effective nuclear charge a net positive charge that is attracting a particular electron
decreases Atomic radius _ across a period (left to right)
paramagnetism something which is _ will be attracted to a magnetic field
diamagnetism something which is _ will be slightly repelled by a magnetic field
isoelectronic same number of electrons
larger traversing down the periodic table, atoms get larger or smaller?
ionization energy the minimum energy needed to remove an electron from an atom or ion in the gas phase
smaller IE _ down the group
larger IE _ across the period to the right