Chapter 2: 2.2 Periodic Trends Flashcards
Define:
Effective nuclear charge
The attractive force felt by a valence electron to the nucleus
or
The net positive charge experienced by the valence electrons when they are shielded (screened) by the core electrons from the full charge of the nucleus
State the trend of the effective nuclear charge as we move across a row of the periodic table from left to right
The effective nuclear charge increases
How does one calculate effective nuclear charge?
Nuclear charge - Core electrons
What is nuclear charge?
Nuclear charge is equal to atomic number
Define:
Atomic radius
Half the length of a bond formed by two atoms of the same element
Define:
Internuclear distance
Distance between 2 nuclei
What are atomic and ionic radii usually expressed in?
Picometers
Angstrom units (100 picometers)
State the relationship between the atomic size and the effective nuclear charge
Inverse relationship
As atomic size decreases, effective nuclear charge increases due to increased valence electron attraction
State the change in size of atoms as one moves down a group
Atomic size increases as more electrons are present, requiring more orbitals with higher n values.
The size of orbitals increase with n and thus with it, the overall size of the atom
What are the exceptions to the general rules of atomic size?
d and f orbitals have poor shielding properties, allowing p electrons to be strongly attracted to the nucleus, making the atomic radius smaller
Why are double and triple bonds typically shorter than single bond lengths?
They contain more electrons that pull the nuclei together
State the trend of bond lengths as one goes across from left to right on the periodic table
Bonds lengths decreases
True or False:
Anions are smaller than neutral atoms
False, anions are larger than neutral atoms and cations are smaller than neutral atoms
Dications are _______ than monocations of the same element
Smaller
Dianions are ______ than monoanions of the same element
Larger
Define:
Ionization energy (IE)
The Minimum amount of energy required to remove a valence electron from a gas-phase atom
Since energy must be ________ to remove an electron, IEs are always ________ quantities
Absorbed
Positive
State the trend for energy required in successive ionization
The energy required for each successive ionization increases because the net positive charge of the nucleus felt by the remaining electrons is increasing. As a result, the remaining electrons are held more tightly to the nucleus
State the correlation between IE and Effective nuclear charge
IE increases from left to right across each row as the effective nuclear charge increases(directly proportional)
Why does the IE decrease when going down a group?
Atoms get larger as we move down a group, the larger distance between the nucleus and the valence electrons means its easier to rip electrons off
Define:
Electron affinity
The change in energy (or enthalpy) in gas phase process
or
The change in energy to gain electron
What are the possibilities for electron affinity?
Thermodynamically favoured (H<0)
Thermodynamically not favoured (H>0)
What are some anomalies for electron affinity?
Group 2
Group 15
Noble gases
Does electron affinity decrease when going down a group?
Yes
Define:
Electronegativity
“The ability of an atom to draw a pair of electrons of a bond to itself”
or
The ability of an atom participating in a chemical bond to draw electron density to itself
State electronegativity trends on the periodic table
Electronegativity tends to decrease going down the groups (due to increasing atomic size)
Electronegativity tends to increase going across the rows
Where are the most electromagnetic elements located on the periodic table?
Top right corner
