CHM 311 Inorganic Vocabulary

enthalpy

A measure of the total heat content of a thermodynamic system

enthalpy change

the amount of heat absorbed or released by a system

standard enthalpy change

the amount of heat absorbed or released by a system in which all products and reactants are in their standard states

exothermic

The content of the products has a lower value than that of the reactants, and heat is released to its surroundings

endothermic

Heat is absorbed from the surroundings by the products

standard heat of formation

The standard enthalpy change for any reaction

catalyst

A substance that causes a reaction to proceed more rapidly to equilibrium

electronic structure

the number and the distribution of electrons about a central nucleus

lines, spectroscopic emissions

The atomic emission spectrums that occur when excited atoms, molecules, or ions release energy as characteristic photons of light or radiation when their electrons drop from higher energy levels to lower energy levels

wave equation

A fundamental mathematical formula in quantum mechanics that describes how the quantum state of a physical system changes over time

wave mechanics

A branch of physics that describes the wave-like behavior of tiny particles like electrons and atoms

orbital

A mathematical region around an atomic nucleus where there is a high probability of finding an electron; a term borrowed from the Bohr concept of orbits

enclosure surface

The depiction of an orbital inside which a majority of the electron density is known to reside

nodes

A point, surface, or plane in an atomic or molecular orbital where the probability of finding an electron is zero

Aufbau principle

The electron configurations are built up from the bottom, using the lowest energy orbitals first

Hund’s rule

When orbitals are available in degenerate sets, maximum spin multiplicity is preserved; that is, electrons are not paired until each orbital in a degenerate set has been half filled

Pauli Exclusion principle

No two electrons may have the same set of four quantum numbers. Where two electrons occupy the same orbital, they must have the opposite spins. An orbital can house at most two electrons.

K shell

The innermost electron shell of an atom, only containing the orbital 1s

L shellT

The second electron energy level in an atomic orbital; n = 2

M shell

The third electron shell surrounding an atomic nucleus; n = 3, and it can contain 18 electrons

noble gases

All gases with outer ns²np⁶ configurations to fall in a column; group 18

alkali metals

The group 1 atoms containing a single s electron in the valence shell

alkaline earth metals

The group 2 atoms containing ns² configuration

halogens

The group 17 elements containing ns²np⁵ electron configurations

chalcogens

The group 16 elements containing ns²np⁴ electron configurations

pnictogens

The elements of group 15 containing ns²np³ electron configurations

main group elements

These elements, which involve outer shells consisting solely of s and p electrons

s-block elements

Chemical elements that have their outermost valence electron in the s-orbital

p-block elements

Chemical elements that have their outermost valence electron in the p-orbital

d-block elements

A group of 40 metal elements found in groups 3 through 12 in the middle section of the periodic table

transition elements

Chemical elements found in the middle block (groups 3 through 12) of the periodic table.

first transition series

The set of elements with incomplete 3d shells

second transition series

The set of elements with partial 4d shells

third transition series

The set of elements with partial 5d shells

lanthanides

A group of 14 metallic chemical elements with atomic numbers from 57 to 71, starting with lanthanum and ending with lutetium. They contain partial 4f orbitals

actinides

A group of 14 metallic chemical elements with atomic numbers 89 , actinium, to 103, lawrencium, where these atoms contain partial 5f orbitals

f-block elements

The elements in the lanthanides and actinides containing partial f orbitals

maximum spin multiplicity

The most stable electronic state that has the largest number of unpaired electron spins

effective nuclear charge (Z_eff)

The net positive charge experienced by an electron in a multi-electron atom after accounting for the repulsive effect of inner-shell electrons

differentiating electron

The final electron added to an atom’s electron configuration when building up elements by atomic number

first ionization enthalpies

The energy required to remove one mole of the outermost electrons from one mole of neutral gaseous atoms to form one mole of gaseous ions with a 1+ charge

single-bond covalent radius

The typical contribution by that element ot the length of a predominantly covalent bond

van der Waals radii

The value that is obtained from the non-bonded distance of closest approach between atoms that are in contact with, but not bonded to one another

ionic radii

A value that represents radii that are assigned to the various ions of the elements as they are more predominantly ionic compounds

electronegativity

An empirical measure of the tendency of an atom in a molecule to attract electrons

paramagnetic

Any atom, ion, or molecule that has one or more unpaired electrons

diamagnetic

Substances that do not contain unpaired electrons

magnetic susceptibility

A measure of the force exerted by the magnetic field on a unit mass of the specimen

magnetic moment

the measured magnetic susceptibility per mole

bonding electrons

The sharing of a pair of electrons between the atoms. These electrons are fixed between the two atoms in Lewis’s definition.

lone-pair electrons

Electrons that are not shared between atoms in Lewis’s definition

saturated systems

The number of valence electrons is just sufficient to provide an octet for each non-hydrogen atom

unsaturated systems

Substances where the number of valence electrons that are available within a molecule or complex ion is not sufficient to allow the Lewis diagram to be written using single-bonds only

multiple bonds

A sharing of more than one pair of electrons in the localized bond approach

electron deficient

Molecules where all of the available valence electrons are used before an octet is achieved for each non-hydrogen atom

octet

A full valence shell

valence shell expansion

Larger atoms may acquire more than an octet of electrons. This is made possible by the availability of valence level d orbitals on these atoms

ligand

The groups that are bonded to the central metal atom

coordinate-covalent bonds

Bonds where both electrons of the metal-ligand bond are supplied by the ligand

Lewis base

An electron-pair donor

Lewis acid

An electron-pair acceptor

18-electron rule

The octet rule is replaced by this rule because of the additional 10 electrons of the d orbitals in any transition series

resonance

The overall result of delocalization of electrons within the structure through the recognition that other Lewis diagrams may be equally valid

hybridization theory

The construction of new orbitals on atoms so that the bonding in a molecule is made to be consistent with its known geometry

VSPER theory

A theory that allows the best geometry for a molecule or polyatomic ion to be predicted

VSEPR

Valence shell electron pair repulsion

sp hybrid

Linear molecule formed

sp² hybrid

Trigonal planar molecule formed

sp³ hybrid

Tetrahedral molecule formed

d²sp³

Octahedral hybridized molecule formed

dsp²

square planar hybridization

sd³

Tetrahedral hybridization

dsp³

Trigonal bipyramidal hybridization or square pyramidal hybridization

occupancy

The sum (x + y) of the AB_xE_y structure. The amount of space around atom A that contains other atoms or lone pairs

isostructural

A molecule having the same structures

isoelectric

A molecule having the same electron configuration

axial

Groups that are positioned above and below the triangular plane in a trigonal bipyramid

equatorial

Groups that are positioned in the triangular plane of a trigonal bipyramid

bond order

The number of electron pairs in bonding molecular orbitals (n_b) and subtract the number of pairs in anti-bonding molecular orbitals (n_a)

group orbitals

The best total, positive overlap with the atomic p_z of the central atom

Important points to remember about the four sigma molecular orbitals

1. In each bonding MO, electron density is large and continuous between adjacent atoms, while in the antibonding MO’s there is a node between each adjacent pair of nuclei.

   1. In each bonding MO, the wave function indicates that an electron pair occupying it is spread out over the whole molecule and is shared by all of the atoms, not just a particular adjacent pair. In other words, in MO’s the electrons are delocalized over the whole extend of the MO.

antibonding

A negative overlap, shared electron density is reduced and internuclear repulsion increases. This creates a net repulsive interaction

Nonbonding

Net overlap is zero and there is neither an increase nor a decrease in shared electron density

Bonding

Overlap has a positive sign when the superimposed regions of the two orbitals have the same sign

MO theory description of a chemical bond

A chemical bond can exist when outer orbitals on different atoms overlap so as to concentrate electron density between the atomic cores

van der Waals forces

All the attractive and repulsive forces that are neither ionic nor covalent. London dispersion forces, dipole-dipole interactions, and Hydrogen bonds

AB₄E₂

Planar structure because of the two lone pairs on the central atom

AB₅E₁

Square pyramidal geometry

AB₆

octahedral geometry, d²sp³ hybridization

AB₂E₃

Linear geometry

AB₃E₂

Distorted planar T shape

AB₄E

Trigonal bipyramid structure with lone pair

AB₅

Perfect trigonal bipyramidal geometry

ABE₃

Unknown geometry

AB₂E₂

Bent geometry, two lone pairs on the sp³ hybridized orbital

AB₃E

Pyramidal and serve as Lewis bases

ML₄ⁿ⁺

Square planar geometry common for d⁸

AO₄^n-

Tetrahedral central atom PO₄^3- and SO₄^2-

AB₄

Always tetrahedral geometry

ABE₂

Linear geometry