Electron configuration and chemical periodicity Flashcards
In 1869, * independently proposed the
*:
• When the * are arranged in order of *, certain sets of properties recur periodically
Dmitri Mendeleev and Lothar Meyer
periodic law
elements, increasing atomic mass,
In *, Dmitri Mendeleev and Lothar Meyer independently proposed the
periodic law:
• When the elements are arranged in order of increasing atomic mass, certain sets
of * recur * .
1869
properties, periodically
Splitting of Energy Levels
• The * nuclear charge Z * orbital energy by * nucleus-electron attractions.
higher, lowers, increasing
Splitting of Energy Levels
• The higher * lowers * by increasing *
nuclear charge Z, orbital energy, nucleus-electron attractions.
An additional electron raises the * due to *.
• Electrons in outer orbitals (higher n) are higher in energy because inner electrons * them from * (effective nuclear charge, Zeff)
orbital energy, electron-electron repulsions
shield, nuclear charge
An additional electron * the orbital energy due to electron-electron repulsions.
• Electrons in * (higher n) are higher in energy because * shield them from nuclear charge (*)
raises
outer orbitals
inner electrons
effective nuclear charge, Zeff
Electrons that have a * probability distribution near the nucleus (*) have * energy. Thus, an * (shell) is split into * (subshell) energies: s < p < d < f .
finite, penetration, lower
energy level, sublevel
Electrons that have a * (penetration) have lower energy. Thus, an energy level () is split into
sublevel () energies: s < p < d < f .
finite probability distribution near the nucleus
shell, subshell
The * of an atom is a designation of how electrons are
* among various orbitals in principal shells and subshells.
electron configuration, distributed
The electron configuration of an atom is a * of how electrons are distributed among various * in *
designation
orbitals
principal shells and subshells
Electrons occupy * in a way that * the energy of the atom.
orbitals, minimizes
Pauli exclusion principle. No two electrons can have * set of * . An orbital can hold a maximum of * electrons, and they must have * spins.
the same, four quantum numbers (n, l,ml ,ms)
two, opposite
Hund’s rule. For orbitals of identical energy (*), electrons initially occupy these orbitals *.
degenerate orbitals, singly (before pairing up)
The Aufbau Principle: electrons fill orbitals starting at the * before filling * (e.g. 1s before 2s).
lowest available (possible) energy states, higher states
The * states that electrons fill orbitals starting at the lowest available (possible) energy states before filling higher states (e.g. 1s before 2s). It is also called *
Aufbau Principle
Building-Up Principle
The order in which these orbitals are filled is given by the * rule, which states
that given two orbitals, the one which has a higher n value but a lower (n + l) value
is considered to be of a * .
• Also called the * rule or the Klechkowski rule.
(n + l)
lower energy level
Madelung
The order in which these orbitals are filled is given by the (n + l) rule, which states
that given two orbitals, the one which has a * value but a * value
is considered to be of a lower energy level.
• Also called the Madelung rule or the * rule.
higher n, lower (n + l)
Klechkowski
Hund’s Rule states that for degenerate orbitals, the * is attained when the number of electrons with the * spin is maximized.
also called Hund’s Rule of *.
lowest energy, same
Maximum Multiplicity
The * refers to the periodic recurrence of certain * when the elements are considered in terms of *
periodic law
physical and chemical properties
increasing atomic number.
The periodic law refers to the * of certain physical and chemical properties when the * are considered in terms of increasing atomic number.
periodic recurrence, elements
The * is an arrangement of the elements, by *, in which elements with similar physical and chemical properties are grouped together
in *
periodic table, atomic number
vertical columns
The periodic table is an * of the elements, by atomic number, in which elements with * are grouped together
in vertical columns
arrangement,
similar physical and chemical properties
A * is a vertical column of elements in the periodic table. Members of a group have *.
group, similar properties
A * is a horizontal row of the periodic table. All members of a period have atoms with *
period
the same highest principal quantum number.
The first two groups–the s block–and the last six groups–the p block–together
constitute the * elements.
Because they come between the s block and the p block, the d block elements
are known as the *.
main-group
transition elements
The first two groups–the * –and the last six groups–the *–together
constitute the main-group elements.
Because they come between the s block and the p block, the * elements
are known as the transition elements.
s block, p block
d block
The f block elements, sometimes called the *
The 15 elements following barium (Z = 56) are called the *, and the
15 following radon (Z = 88) are called the *
inner transition elements,
lanthanides
actinides
The * elements, sometimes called the inner transition elements,
The 15 elements following * are called the lanthanides, and the 15 following * are called the actinides
f block
barium (Z = 56)
radon (Z = 88)
The atomic radius of an element is a measure of the *, usually the distance from the * to the *
size of the atom
nucleus, boundary of the surrounding cloud of electrons
- is one-half the distance between the centers of two atoms that are bonded covalently. It is the atomic radius associated with an element in its
covalent compounds. - is the radius of a spherical ion. It is the atomic radius associated with an element in its ionic compounds.
- is one-half the distance between the centers of adjacent atoms
in a solid metal.
Covalent radius
Ionic radius
Metallic radius
The * are strictly hard sphere radii measured using atomic distances in closest packed *.
• solid sample of a noble gas
van der Waals radii,
crystals
The atomic radius tends to * as one progresses across a period from left to right
• because the effective nuclear charge (Zeff) *, thereby * the orbiting electrons and lessening the radius
decrease
increases, attracting
The * tends to decrease as one progresses across a period from left to right
• because the * (Zeff) increases, thereby attracting the
orbiting electrons and * the radius
atomic radius, effective nuclear charge
lessening
The atomic radius usually * while going down a group
• due to the addition of a new *
increases energy level (shell).
Cations are * than the atoms from which they are formed.
• For isoelectronic cations, the more * the ionic charge, the * the ionic radius.
smaller
positive, smaller
Anions are * than the atoms from which they are formed.
• For isoelectronic anions, the more * the charge, the * the ionic radius.
larger
negative, larger
The ionization energy, I, is the quantity of energy a gaseous atom must * to
be able to * an electron
absorb, expel
The * is the energy required to remove the * electron from a * atom.
first ionization energy I1
most loosely held, gaseous
The second ionization energy I2 is the energy required to remove an electron
from a *.
gaseous unipositive ion
Ionization energies * as atomic radii *.
• IE * down a group.
• IE * across a period.
decrease, increase
decreases, increases
Electron affinity is the energy * associated with the * of an electron by a * atom.
change, gain, neutral gaseous
The smaller atoms to the right of the periodic table (e.g., group 17) tend to have
* electron affinities.
• EAs tend to become * in progressing toward the bottom of a group, with the notable exception of the second-period members of groups 15, 16, and
17 (namely, N, O, and F).
• Some atoms have no tendency to gain an electron, such as the noble gases where
electrons have to enter the next shell, and groups * where the electrons
have to enter an empty p subshell, etc.
large, negative
less negative
2 and 12
Electronegativity (EN) is a measure of the * of an atom to * electrons towards itself in the context of a chemical bond.
ability, attract
EN is related to IE and EA. In 1934, * developed an approach
where
EN = *
Robert S. Mulliken
IE − EA2
As one moves from left to right across a period in the periodic table, the
electronegativity *
• due to the stronger attraction that the atoms obtain as the * increases
increases, nuclear charge
Moving bf down a group, the electronegativity *
• due to the longer * between the nucleus and the valence electron shell,
thereby * the attraction.
decreases
distance
decreasing
In general, EN is * related to atomic size.
inversely
An important property related to the electron configurations of atoms and ions
is their * in a magnetic field. A spinning electron is an electric charge in
*, which induces a *
behavior
motion, magnetic field
• A diamagnetic substance has all its electrons * and the individual magnetic effects cancel out. It is slightly repelled by a magnetic field.
• A paramagnetic substance has one or more * electrons in its atoms or molecules and the individual magnetic effects do not cancel out. It is attracted
into a magnetic field.
paired
unpaired
A metal is an element whose atoms have * numbers of electrons in the * electronic shell. Removal of an electron(s) from a metal atom occurs without great difficulty, producing *.
small, outermost
a positive ion (cation)
Metals generally have a * appearance, are *, and are
able to *.
lustrous
malleable and ductile
conduct heat and electricity
A nonmetal is an element whose atoms tend to gain small numbers of electrons to form * with the electron configuration of a *.
Nonmetals are mostly *, *(bromine), or * solids and
are very * conductors of heat and electricity.
negative ions (anions), noble gas
gases, liquid, low melting point
poor
A * is an element that may display both metallic and nonmetallic properties under the appropriate conditions.
metalloid
Metallic property * across a period with * in number of valence electrons, as well as * in atomic radius, and it * down the group
with increase in the *
decreases
increase, decrease
increases
number of shells and atomic radius.
