Chapter 08: Periodic Properties of the Elements Flashcards
Periodic law
When elements are arranged in order of increasing mass, certain properties recur periodically
Electron configuration
Shorthand for showing particular orbitals that electrons occupy for that atom
e.g. 1s1 for Hydrogen
Pauli exclusion principle
No two electrons in an atom can have the same four quantum numbers
Each orbital can have a max of 2 electrons, with opposing spins
Degenerate
Having the same energy
Coulomb’s law
Like charges:
Potential energy is positive
Decreases as particles are farther apart
Repel each other
Opposite charges:
Potential energy is negative
Becomes more negative as particles are closer
Opposites attract
Magnitude of interaction increases as charges of particles increase
i.e. 1- and 2+ are more strongly attracted than 1- and 1+
Shielding
Screening or repulsion of one electron by other electrons
Effective nuclear charge
Zeff = Z - S
= actual nuclear charge (atomic #) - charge screened (by core electrons)
Increases as you go right on periodic table
(Which is why radius gets smaller to the right)
Penetration
When an outer electron penetrates into lower energy region, it experiences a greater nuclear charge, and thus a lower energy
Orbital diagram
A diagram that represents the orbitals occupied by electrons in a given atom
Boxes for arrows of electrons, with orbital name under each box
e.g. 1s with a box above
Aufbau principle
Lower energy orbitals fill before higher energy orbitals to minimize the energy of the atom
Hund’s rule
Electrons first occuply orbitals of equal energy singly with parallel spins
Transition elements with irregular electron configurations (10)
- Cr (Chromium) 24
- Cu (Copper) 29
- Nb (Niobium) 41
- Mo (Molybdenur) 42
- Ru (Rutherium) 44
- Rh (Rhodium) 45
- Pd (Palladium) 46
- Ag (Silver) 47
- Pt (Platinum) 78
- Au (Gold) 79
Paramagnetic
Electron configurations with unpaired electrons
Have a net magnetic field
Attracted to a magnetic field
Diamagnetic
All electrons are paired in electron configuration
No magnetic field of its own
Slightly repelled by a magnetic field
Isoelectronic
Elements that have the same number of electrons & same ground-state electron configuration
Van der Waals radius
Nonbonding radius of an atom
Covalent radius
Radius of an atom when it’s bonded
Atomic radius
Average radius of a bonded atom based on measuring large numbers of elements and compounds
Periodic trends in atomic radius for transition metals
Remains roughly constant across the d block
Ionic radius
Cation is smaller than parent atom (e.g. Li+ v. Li)
Anion is always larger than parent atom (e.g. Cl- v. Cl)
Cation v. anion radius size
Cations < anions
Except Rb+ and Cs+ are larger than F- and O2-
Ionic radii of isoelectronic species
Larger positive = smaller cation
Larger negative = larger anion
Ionization energy
IE1 + X (g) → X+ (g) + e-
The minimum energy required to remove an electron from a gaseous atom or ion
IE1 < IE2 < IE3
(Each subsequent electron removal requires more energy, due to increasing positive charge of atom)
(Large increase in IE when core electrons are removed)
Periodic trends in first ionization energy
The larger the Zeff, the more energy it takes to remove it.
*IE1 increases to the right
**Except: Group 2A to 3A & Group 5A to 6A
The farther the electron is from the nucleus, the less energy it takes to remove it
*IE1 decreases toward the bottom
Electron affinity
X (g) + e- → X- (g) + EA
The energy released (-kJ/mol) when a neutral gaseous atom gains an electron
The more energy released when electron is gained, the more negative the EA (-kJ/mol)
Periodic trends in electron affinity (6)
- EA becomes less negative going down Group 1A
- 3rd Period becomes more negative from 2nd Period
- EA becomes more negative going right
- Group 5A has less negative EA than expected (extra electron pairs up)
- Groups 2A/8A have less negative EA (extra electron goes into new level/sublevel)
- Group 7A has most negative EA
Periodic trends in metallic character
Metallic character increases going left and down
Periodic trends in Group 1A (alkali metals) (7)
- IE increases up the group (as with most groups)
- Generally low IEs (want to lose electron)
- EA becomes more negative down the group
- Melting point increases up the group (unusual)
- Generally low melting points
- Density increases down the group (except K)
- Reactivity increases down the group
Periodic trends in Group 7A (halogens) (4)
- IE increases up the group (as with most groups)
- EAs are high (want to gain electrons)
- React with H to form binary acids
- Mass increases more than volume
Periodic trends in Group 8A (noble gases) (3)
- IE increases up the group (as with most groups)
- IEs are high (don’t want to lose electrons)
- Only Kr & Xe form compounds, usually with F
Metals in Group 1A vs. 1B (2)
- Similar outer electron configurations: ns1
- Properties different because different IEs
1A: lower IE, more reactive (want to lose ns1)
1B: higher IE, less reactive
Diagonal relationship
- Li & Mg
- Be & Al (Al is a metal like 2A)
- B & Si (metalloids)
Similar because of charge density (ion charge / volume)
Properties of oxides across a period (3)
- Metals/Groups 1A/2A form basic oxides
- Nonmetals form acidic oxides
- Al form amphoteric oxides (both basic/acidic)
