Chapter 7 Flashcards
S orbital
1 orbital
Circular cloud
P orbital
3 orbitals
Figure 8
D orbital
5 orbitals
2 figure 8’s
F orbital
7 orbitals
Complex form of a figure 8
What makes an atom more paramagnetic?
It has more unpaired electrons
Coulombic force of attraction
The force of attraction between the negative electrons in the shells, and the positive protons in the nucleus
Effective nuclear charge
The charge that of the nucleus that an electron actually experiences
Shielding electrons
Block the outer electrons of an atom from feeling the full nuclear charge
Equation for effective nuclear charge:
Zeff= atomic # - # of shielding electrons
Atomic radius trend
- Decreases across a group because atoms gain protons and electrons are pulled closer into the nucleus
- Increase down a group because the number of shells increase
First ionization energy definition
The minimum amount of energy needed to remove one electron from an atom’s valence shell
First ionization energy trend
- Increases across a period because there are more valence electrons
- Decreases down a group because there are more shells and as the valence electrons get farther from the nucleus, they are easier to remove
First ionization energy anomalies
Groups 2 and 3: Group 3 atoms have a lower first ionization energy than group 2 because their electrons are in the p orbital and therefore slightly farther from the nucleus
Groups 15 and 16: Group 15 has all unpaired electrons so they are slightly harder to remove since there are less repelling forces in the atom
Electronegativity definition
A measure of how much an atom attracts electrons to bind with it
Electronegativity trend
- Increases across a period because there are more valence electrons that want to gain an electron
- Decrease down a group because the electrons get farther from the nucleus so it’s harder for the nucleus to pull in electrons
Electron affinity defintion
The likelihood of an atom to gain an electron
Electron affinity trend
- Increases across a period because there are more valence electrons that want to gain an electron
- Decrease down a group because the electrons get farther from the nucleus so it’s harder for the nucleus to pull in electrons
Metal reactivity
Metals are typically oxidized (lose electrons)
Metal reactivity trend
- Decreases across a period because there are more valence electrons making them harder to remove
- Increases down a group because there is more distance from the nucleus and therefore the electrons are easier to remove
Non-metal reactivity
Non-metals are typically reduced (gain electrons)
Non-metal reactivity trend
- Increases across a period because there are more valence electrons that are closer to having full shells
- Decreases down a group because as the electrons get farther from the nucleus, the nucleus has less of a pull
Isoelectronic definition
Particles with the same number of electrons
Hund’s rule
A single electron must be placed into each orbital before any pairing takes place
Pauli exclusion principal
No more than 2 electrons in each orbital, electrons must have opposite spins
