Chapter 7/8: Principles of Atomic Theory, Orbitals, Quantum-Mechanical Model of the Atom Flashcards
Atomic radius increases as you move…
DOWN a group in the periodic table
generic outer electron configuration for Halogens
ns2np5
generic outer electron configuration for Noble Gases
ns2np6
the release of energy corresponds to a transition to a more stable state
the element that is most likely to gain an electron should be placed at the top of the ranking
elements with a full s or p subshell are particularly stable, so the addition of an electron…
is unfavorable/less likely (i.e. Fl will gain 1 electron (1-) sooner than Al, which will gain ?5? (5-) sooner than Be, which will give up 2 (2+)
The matter-wave of the electron occupies the space near the nucleus and is continuously influence by it;
*The Schrodinger wave equation allows us to solve for the energy states associated with a particular atomic orbital
The square of the wave function gives the probability density, a measure of the probability of finding an electron of a particular energy in a particular region of the atom
Quantum-Mechanical model of atom
- The principle quantum number (n) is a positive integer; The value of n indicates the relative size of the orbital and therefore its relative distance from the nucleus
- The angular momentum quantum number (l) is an integer from 0 to (n-1); The value of (l) indicates the shape of the orbital
- The magnetic quantum number (m sub l) is an integer with values from -(l) to +(l); The value of m sub l indicates the spatial orientation of the orbital
Quantum numbers and atomic orbitals:
an atomic orbital is specified by 3 Quantum Numbers: n, l, m sub l
(and also m sub s)
Each electron is described completely by all 4 Q #s; the first 3 describe the orbital, and the last describes the electron spin direction (up or down)
No two electrons in the same atom can have the same 4 quantum #s
An atomic orbital can hold a maximum of 2 electrons and they must have opposing spins
Pauli’s exclusion principle
*Nuclear charge (Z) and shielding by other electrons
Greater nuclear charge increases nucleus-electron interactions and lowers sublevel energy
-shielding by other electrons reduces full nuclear charge to an effective nuclear charge (Z eff)–the nuclear charge an electron actually experiences
-orbital shape also effects sublevel energy
Factors affecting orbital energies
- Electrons in the same energy level shield each other to some extent
- Electrons in inner energy levels shield outer electrons very effectively–the further from the nucleus an electron is, the lower its (Z eff)
Shielding and Orbital Energy
Orbital shape causes electrons
in some orbitals to “penetrate”
close to the nucleus.
Penetration increases nuclear
attraction and decreases
shielding.
Penetration and sublevel energy.
Order of sublevel energies:
s < p < d < f
For a given n value, a lower l value indicates a lower energy sublevel.
Each energy level is split into sublevels of differing energy.
Splitting is caused by penetration and its effect on shielding.
electrons are always placed in the lowest sublevel energy available
the Aufbau principle
when orbitals of equal energy are available, the
lowest energy electron configuration has the maximum number of unpaired
electrons with parallel spins.
Hund’s Rule
Elements in the same group of the periodic table have the same outer electron
configuration.
Elements in the same group of the periodic table exhibit similar chemical
behavior.
Similar outer electron configurations correlate with similar chemical
behavior.
Electron Configuration and Group
Atomic size increases as the principal quantum number n increases.
As n increases, the probability that the outer electrons will be further from the nucleus increases.
Atomic size decreases as the effective nuclear charge Zeff increases.
As Zeff increases, the outer electrons are pulled closer to the nucleus.
Trends in atomic size
For main group elements:
atomic size increases down a group in the periodic table and decreases across a period.
Ionization energy (IE) is the energy required for the complete removal of 1 mol of electrons from 1 mol of gaseous atoms or ions.
Ionization energy tends to decrease down a group and increase across a period.
Atoms with a low IE tend to form cations.
Atoms with a high IE tend to form anions (except the noble gases).
Trends in Ionization Energy
Electron Affinity (EA) is the energy change that occurs when 1 mol of electrons is added to 1 mol of gaseous atoms or ions.
Atoms with a low EA tend to form cations.
Atoms with a high EA tend to form anions.
The trends in electron affinity are not as regular as those for atomic size or IE.
Trends in Electron Affinity
Reactive nonmetals have high IEs and highly negative EAs.
These elements attract electrons strongly and tend to form negative ions in ionic compounds.
Behavior patterns
Reactive metals have low IEs and slightly negative EAs.
These elements lose electrons easily and tend to form positive ions in ionic compounds.
Behavior patterns
Noble gases have very high IEs and slightly positive EAs.
These elements tend to neither lose nor gain electrons.
Behavior patterns
Atomic size: increases LEFT and DOWN
Ionization energy: increases UP and RIGHT
Electron Affinity: loose trend UP and RIGHT (many exceptions)
Trends in the 3 atomic properties
*typically shiny solids with moderate to high melting points.
•good conductors of heat and electricity, and can easily be shaped/malleable
•tend to lose electrons and form cations (they are easily oxidized.)
•generally strong reducing agents.
*most form ionic oxides, which are basic in
aqueous solution.
Metallic Behavior
A pseudo-noble gas configuration is attained when a metal atom empties its highest energy level.
The ion attains the stability of empty ns and np sublevels and a filled (n-1)d sublevel.
Sn: ([Kr]5s2 4d10 5p2) → 4(e-) + Sn4+([Kr]4d10)
A metal may lose only the np electrons to attain an inert pair configuration.
The ion attains the stability of a filled ns and (n-1)d sublevels.
Sn:([Kr]5s2 4d 10 5p2)→ 2(e-)+ Sn2+([Kr]5s2 4d10)
Electron configurations of mono-atomic ions
A species with one or more unpaired electrons exhibits paramagnetism – it is attracted by a magnetic field.
A species with all its electrons paired exhibits diamagnetism – it is not attracted (and is slightly repelled) by a magnetic field.
Magnetic properties of Transition Metal Ions
Cations are smaller than their parent atoms
Anions are larger than their parent atoms
Ionic radius increases down a group as n increases
Cation size decreases as charge increases
An isoelectronic series is a series of ions that have the same electron configuration. Within the series, ion size decreases with increasing nuclear
charge.
Ionic size v. Atomic size
