Chapter 3: Periodic Properties of the Elements Flashcards
Who’s insight led to the development of the periodic table?
Dmitri Mendeleev
Describe the density trend on the periodic table.
The density of elements tends to increase as we move down a column in the periodic table. This is true because the mass of each successive atom increases even more than its volume does.
Define periodic properties.
A property of an element that is predictable based on an element’s position in the periodic table.
Define periodic law.
The law based on the observation that when the elements are arranged in order of increasing mass, certain sets of properties recur periodically.
Who explained why elements were ordered by atomic number?
Henry Moseley
Define main-group elements.
One of the elements found in the s or p block of the periodic table, whose properties tend to be predictable based on their positions in the table.
Define transition elements (transition metals).
One of the elements found in the d block of the periodic table whose properties tend to be less predictable based simply on their position in the table.
Define family (group).
On the periodic table, one of the columns within the main group elements; a family or group of elements exhibits similar chemical properties.
Define electron configuration.
A notation that shows the particular orbitals that are occupied by electrons in an atom.
Define ground state.
The lowest energy state of an atom or molecule.
Define orbital diagram.
A diagram similar to an electron configuration that symbolizes an electron as an arrow in a box representing an orbital, with the arrow’s direction denoting the electron’s spin.
Define Pauli exclusion principle.
The principle stating that no two electrons in an atom can have the same four quantum numbers.
Define degenerate.
Describes two or more electron orbitals with the same value of n that have the same energy.
Describe the energy between s, p, d, and f orbitals.
E(s orbital) < E(p orbital) < E(d orbital) < E(f orbital)
Define Coulomb’s law.
The law that states that the potential energy (E) of two charged particles depends on their charges (q>1 and q>2) and on their separation (r).
Formula Coulomb’s Law
E = (1/4pi e0)(q1q2 / r)
Conclusions from Coulomb’s Law
- PE from like charges is positive, but decreases as the particles get further apart. Like to move towards lower PE, therefore like charges repel one another.
- PE from unlike charges is negative and becomes more negative as the particles get closer together. Unlike charges attract one another.
- The magnitude of interaction increases as the charges of the particles increase. An Electron with a charge of 1- is more attracted to a nucleus with 2+ charge than 1+ charge.
Define shielding.
The effect on an electron of repulsion by electrons in lower-energy orbitals that screen it from the full effects of nuclear charge.
Define effective nuclear charge (Z>eff).
The actual nuclear charge experienced by an electron, defined as the charge of the nucleus plus the charge of the shielding electrons.
Define penetration.
The phenomenon in which some higher-level atomic orbitals have significant amounts of probability with the space occupied by orbitals of lower energy level. For example, the 2s orbital penetrates into the 1s orbital.
Define aufbau principle.
The principle that indicates the pattern of orbital filling in an atom.
Define Hund’s rule.
The principle stating that when electrons fill degenerate orbitals, they first fill them singly with parallel spins.
Define valence electrons.
The electrons that are important in chemical bonding. For main-group elements, the valence electrons are those in the outermost principal energy level. For transition elements, we also count the outermost d electrons among the valence electrons.
Why do elements in a column of the periodic table have similar chemical properties?
They have the same number of valence electrons.
