week 7 Flashcards
transition metals and trends (29 cards)
Transition d block elements
where on pt
ox states + what they mean
Elements from group 3 to 13
They exist in low (1+, 2+) and high (3+ to 7+ or higher) oxidation states.
Low oxidation sates = similar to s and p blocks
High oxidation states = coordination chemistry
What is the atomic size trend?
atomic radius increases down a group and decreases left to right across a period
What is coordination chemistry?
the study of transition metal complexes
What does degenerate mean?
describes orbitals with the same energy
What is the electron affinity trend?
Electron affinity becomes less negative down a group. As the principal quantum number increases, the size of the orbital increases and the affinity for the electron is less.
What is the electronegativity trend?
as you move down a group on the periodic table, the electronegativity of an element decreases because the increased number of energy levels puts the outer electrons very far away from the pull of the nucleus.
Electronegativity
increases as you move from left to right across a period on the periodic table
What is the ionisation energy trend?
the general trend is for ionisation energies to increase across a period
What is a radial node?
A node is a point where the electron probability is zero
. Radial nodes are spheres (at fixed radius) that occurs as the principle quantum number increases
What is an oxidation state?
the charge that an atom in a molecule or ion would have if all of the electrons in its bonds belonged entirely to the more electronegative atoms
How many nodal planes do (four of the) d-orbitals have?
Two
What is the electron configuration of Chromium?
[Ar] 4s^1 3d^5
What is the electron configuration of copper?
[Ar] 4s^1 3d^10
True or False: In forming ions, s electrons are lost before d electrons
True
What is the electron configuration of Cu(II)?
[Ar] 3d^9
What is an angular node?
A node is a point where the electron probability is zero. Angular nodes are typically flat plane (at fixed angles), like those in the diagram above. The ℓ quantum number determines the number of angular nodes in an orbital
shielding what? zeff? what orbital type has more?
= describe trends of the transition metals in the periodic table.
Shielding is the cancelling out of a portion of the attraction of the nucleus to the electron, due to the repulsion of these electrons to each other.
shielding effect is related to the effective nuclear charge (Zeff) because Zeff is defined as the nuclear charge that valance electrons experience i.e. the valence electrons are pulled towards the nucleus by Zeff.
shielding of core electrons also depends on the nature of the orbital
s > p > d > f (relative shielding effect of electrons found in different orbitals)
Ionization energy definition
1st v 2nd
Ionisation energy (IE) is the energy required for the complete removal of 1 mol of electrons from 1 mol of atoms or ions in the gaseous state. It can be summarised by the following:
IE1 = first ionisation energy: removes an outermost electron from the gaseous atom:
atom(g) → ion+(g) + e- ∆E = IE1 > 0
IE2 = second ionisation energy: removes a second electron from the gaseous ion:
ion+(g) → ion+2(g) + e- ∆E = IE2 > IE1
atoms with low IE
tend to form cations during reactions
high ionisation energy,
any s-orbital
e- subshells
(or s-orbital) can be empty (i.e. no electrons present in this orbital), half-filled (i.e. when one electron is present) or fully-filled (i.e. when two electrons are present).
any p-orbital
e- subshells
there are three p-orbitals.
half-filled electronic configuration is achieved when three electrons are filled in this subshell according to the Hund’s rule (not engertically favourable)
Electron Affinity def
Electron Affinity
Electron affinity can be described as the energy change accompanying the addition of 1 mol of electrons to 1 mol atoms or ions in a gaseous state
Electronegativity
what?
tasntion metals
Simply put, electronegativity is a measure of an atoms tendency to attract a bonding pair of electrons.
Similar to the ionisation energy, the electronegativity of transition metals increases across the row but then decrease down a group
d-orbitals
They are five-fold degenerate (all equal energy)
and like p-orbitals are directional (Figure 8)
3dxy 3dxz 3dyz 3x2-y2
3dz2
Orbital filling and electronic configuration
Hund’s rule pairing of electrons cannot begin until each orbital in the set contains one electron. Electrons in a degenerate set will singly occupy these orbitals in parallel spins
The Pauli exclusion principle states that no two electrons in an atom can be defined by the same four quantum numbers n, I, ml, and ms
The Aufbau principle states that in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels. For example, the 1s shell is filled before the 2s subshell is occupied
