AP Chemistry Atomic Structure: Atomic Number = # of protons = # of electrons in a neutral atom Mass Number = protons and neutrons (Isotopes) Mass #-238 U Protons=92 Electrons 92 Atomic # 92 Neutrons=238-92=146 Electron Configurations: 1. Order of filling: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 proper way of writing: 1s 2 2s 2 p 6 3s 2 p 6 d 10 4s 2 p 6 d 10 f 14 5s 2 p 6 d 10 f 14 … 2. Atoms gain or lose electrons to obtain a filled octet a) when transition metals lose electrons, they lose from the “s” sublevel first ex. Fe: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 Fe 3+ : 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 b) metals get oxidized (lose electrons) to form cations (pos.) while nonmetals get reduced (gain electrons) to form anions (neg.). 3. Quantum #‟s n = energy level = (1,2,3…) L = sublevel = (0 → n-1) m L = orbital = - L → + L m s = spin # = ± ½ (clockwise or counterclockwise) 4. Oxidation states vs. group # (Group # = highest possible oxidation state) Group: I = +1 IV = +4 (Except Carbide =C -4 ) VII = -1 II = +2 V = -3 VIII = 0 III = +3 VI = -2 Transition metals have multiple oxidation states but (+2) is the most common b/c of “s” sublevel being lost first. *Hund‟s rule: diamagnetic: No unpaired electrons Pauli Exclusion Principle paramagnetic: unpaired electrons Shielding Effects (penetration) ferromagnetic: Fe, Co, Ni Heisenberg Uncertainty Principle
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AP Chemistry
Atomic Structure:
Atomic Number = # of protons = # of electrons in a neutral atom
Mass Number = protons and neutrons (Isotopes)
Mass #-238
U Protons=92 Electrons 92
Atomic # 92
Neutrons=238-92=146
Electron Configurations:
1. Order of filling: 1s22s
22p
63s
23p
64s
23d
104p
65s
24d
105p
66s
24f
145d
106p
67s
2
proper way of writing: 1s22s
2p
63s
2p
6d
104s
2p
6d
10f14
5s2p
6d
10f14
…
2. Atoms gain or lose electrons to obtain a filled octet
a) when transition metals lose electrons, they lose from the “s” sublevel first
ex. Fe: 1s22s
22p
63s
23p
64s
23d
6 Fe
3+ : 1s
22s
22p
63s
23p
63d
5
b) metals get oxidized (lose electrons) to form cations (pos.) while nonmetals get
reduced (gain electrons) to form anions (neg.).
3. Quantum #‟s n = energy level = (1,2,3…)
L = sublevel = (0 → n-1)
mL = orbital = - L → + L
ms = spin # = ± ½ (clockwise or counterclockwise)
4. Oxidation states vs. group # (Group # = highest possible oxidation state)
Group: I = +1 IV = +4 (Except Carbide =C-4
) VII = -1
II = +2 V = -3 VIII = 0
III = +3 VI = -2
Transition metals have multiple oxidation states but (+2) is the most common b/c of “s”
sublevel being lost first.
*Hund‟s rule: diamagnetic: No unpaired electrons
Pauli Exclusion Principle paramagnetic: unpaired electrons
Shielding Effects (penetration) ferromagnetic: Fe, Co, Ni
Heisenberg Uncertainty Principle
Periodic Trends:
I. Atomic Size(radius) : On Periodic Table Increases ←
↓ a) decreases left to right b/c electrons are in same energy level therefore they do not
add size but nuclear charge increases pulling electron cloud in more tightly
b) increases going down in a group b/c adding more energy levels
Greater difference in size between energy levels 1,2, & 3 than between 4,5,6, & 7
b/c energy levels are not evenly spaced.
(Transition metals are all nearly the same size)
II. Ionization Energy = inversely related to size (the bigger the atom, the lower the
ionization energy)
a) exceptions to trend occur when electrons are in filled or ½ filled sublevel
Between groups IIA and IIIA - Increased shielding by "s" electrons
b) Between groups VA and VIA - Increased electron ↔ electron
repulsions because of electrons beginning to pair up in "p" orbitals
c) ions become very stable (high I.E.) once they obtain a noble gas config.
III. Electron Affinity: Opposite of Ionization Energy
Groups IIA, IIB, & VIII have least attraction (most negative value) for extra electrons b/c
they are already stable
Intermolecular Forces 4 types of substances
1. Ionic: metal w/ nonmetal
a) High melting & boiling points b/c of high lattice energy binding ions together