1
Q
What trend occurs in atomic radius across a period from left to right?
A
Atomic radius decreases due to increasing effective nuclear charge (Zeff) pulling valence electrons closer.
2
Q
What trend occurs in atomic radius down a group from top to bottom?
A
Atomic radius increases as electrons occupy higher principal quantum numbers (n), adding shells further from the nucleus.
3
Q
What is an ion?
A
A charged atom.
4
Q
What are cations?
A
Positively charged ions formed by losing electrons.
5
Q
What are anions?
A
Negatively charged ions formed by gaining electrons.
6
Q
How does the size of cations compare to their parent atoms?
A
Cations are smaller than their parent atoms due to increased Zeff and reduced electron-electron repulsion.
7
Q
How does the size of anions compare to their parent atoms?
A
Anions are larger than their parent atoms due to decreased Zeff per electron and increased electron-electron repulsion.
8
Q
What is effective nuclear charge (Zeff)?
A
The net attractive force felt by valence electrons.
9
Q
What is ionization energy (IE)?
A
The energy required to remove an electron from a gaseous atom or ion.
10
Q
What is the trend in ionization energy across a period?
A
Ionization energy increases due to smaller atomic size and higher Zeff.
11
Q
What is the trend in ionization energy down a group?
A
Ionization energy decreases as valence electrons are further from the nucleus and experience more shielding.
12
Q
What happens to successive ionization energies?
A
Removing each additional electron generally requires more energy, with large jumps when removing electrons from stable, closed-shell configurations.
13
Q
What is electron affinity (EA)?
A
The energy change when an electron is added to an atom.
14
Q
Which elements have large negative electron affinities?
A
Halogens, as they release energy when gaining an electron.
15
Q
What are the main groups in the periodic table?
A
Group 1: Alkali metals, Group 2: Alkaline earth metals, Group 16: Chalcogens, Group 17: Halogens, Group 18: Noble gases.
16
Q
What is allotropy?
A
Different structural forms of the same element with distinct properties.
17
Q
Give an example of allotropes of carbon.
A
Diamond and graphite.
18
Q
What is formal charge (FC)?
A
FC = (valence electrons) − (nonbonding electrons) − (number of bonds).
19
Q
What is resonance in chemistry?
A
The concept of delocalized electrons in molecules that can be represented by multiple structures.
20
Q
What types of anions does oxygen form?
A
Oxide (O2−), peroxide (O2^2−), and superoxide (O2^−).
21
Q
What is oxidation?
A
The loss of electrons.
22
Q
What is reduction?
A
The gain of electrons.
23
Q
What is the trend in halogens down the group?
A
Melting/boiling point increases, density increases, atomic radius increases, ionization energy decreases.
24
Q
What are noble gases known for?
A
Having stable electron configurations and being unreactive.
25
What is the octet rule?
Atoms tend to form bonds to achieve a stable octet (eight valence electrons).
26
How do transition metals exhibit variable oxidation states?
They can lose both ns and (n−1)d electrons in bonding.
27
What is the significance of Roman numerals in naming ionic compounds?
They specify the oxidation state of transition metals in compounds.
28
What is the charge of a negative ion from Group 16?
−2.
29
What is the electron configuration of neutral iron (Fe)?
[Ar] 3d6 4s2.
30
What is the electron configuration of copper(I) (Cu+)?
[Ar] 3d10.
31
What is the measured O–O bond length in O3?
~1.278 Å.
32
What happens to bond length as the number of bonds increases?
Bond length decreases.
33
What are the trends in reactivity for superoxides?
Often stronger oxidizing agents than peroxides.
34
What is the charge balance in neutral compounds?
The sum of all oxidation numbers equals 0.
35
What is the trend in boiling points for noble gases down the group?
Boiling point increases.
36
Define the term 'chemical bond'.
The force of attraction which holds atoms together.
37
What is the 'cross-over method' in ionic compound formation?
Using ionic charges to balance stoichiometry to find the empirical formula.
38
What are common charges for nonmetals in Group 15?
−3.
39
What is the electronic configuration of manganese (Mn)?
[Ar] 3d5 4s2.
40
What is the electronic configuration of silver (Ag)?
[Kr] 4d10 5s1.
41
What is the significance of half-filled (d5) or fully filled (d10) d subshells?
They are associated with stability in transition metals.
42
What is the trend in density for heavier chalcogens and halogens?
Density increases down the group.
43
What is the relationship between ionization energy and atomic radius?
Smaller atomic radius typically correlates with higher ionization energy.
44
What is the electron configuration of cobalt (Co)?
1s2 2s2 2p6 3s2 3p6 4s2 3d7.
45
What is a practical example of ionization energy with calcium (Ca)?
Ca → Ca+ + e− (IE1), Ca+ → Ca2+ + e− (IE2), Ca2+ → Ca3+ + e− (IE3, very high).
46
What are the main groups of elements where d orbitals are not involved?
The 'A' groups of main-group elements (s and p blocks).
47
What is the trend in ionization energy for halogens?
They generally have relatively high ionization energy compared to metals.
48
What is the trend in melting and boiling points for nonmetals?
Generally increase down the group.
49
What are redox reactions?
Reactions involving the transfer of electrons between species
50
What is the significance of variable oxidation states?
Foundational for coordination chemistry and catalysis
51
Define lattice energy.
Links ionic bonding strength to charges and radii (Born–Haber cycle)
52
What are periodic anomalies?
Exceptions in periodic trends, explained by subshell stability (e.g., Cr, Cu, Ag)
53
What is the first step in balancing redox reactions?
Assign oxidation states
54
What is the electron configuration practice suggestion?
Write configurations for main-group and transition metals, then derive likely ion charges
55
What is resonance in chemistry?
Delocalized electrons with an average bond order
56
What trend occurs in atomic radius across the periodic table?
Decreases across, increases down
57
What does effective nuclear charge (Zeff) do across the periodic table?
Increases across
58
How does ionization energy change across and down the periodic table?
Increases across, decreases down
59
What is unique about halogens' electron affinity?
Halogens have highly negative electron affinities
60
What happens to the size of cations and anions?
Cations are smaller; anions are larger
61
What is the oxidation state of oxides?
O2−
62
What is the oxidation state of peroxides?
O2^2−
63
What is the oxidation state of superoxides?
O2−
64
What defines halogens?
Diatomic, strong oxidizers, salt-formers
65
What characterizes noble gases?
Monatomic, inert, can form compounds under extreme conditions (e.g., XeF2)
66
What type of bonding occurs between metals and nonmetals?
Ionic bonding following the octet rule
67
What is a key characteristic of transition metals?
Variable oxidation states; lose ns before (n−1)d
68
How are variable-charge metals named?
Using Roman numerals (e.g., iron(II), iron(III))
69
What practice targets should be focused on in chemistry?
* Electron configurations
* Formal charge
* Redox reactions
* Balancing equations
70
Why can noble gases form compounds under extreme conditions?
Due to unique electronic configurations and stability
71
What is a notable property of O2 despite its double bond in Lewis structures?
Paramagnetism explained by MO theory
72
What is an ionic bond?
An electrostatic force of attraction between two oppositely charged ions.
73
What type of ions do metals and non-metals form in ionic bonding?
Metals form cations (positively charged ions) and non-metals form anions (negatively charged ions).
74
What is Ionization Energy (IE)?
The energy required to remove an electron from a gaseous atom or ion, always positive.
75
What is Electron Affinity (EA)?
The energy change that occurs when a gaseous atom gains an electron, typically negative.
76
What is Lattice Energy?
The energy released when gaseous ions come together to form a solid crystalline lattice, always large and negative.
77
What does Coulomb's Law state regarding lattice energy?
The force of attraction between charged particles is directly proportional to the product of the charges and inversely proportional to the distance between them.
78
How does charge affect lattice energy?
Greater charge leads to greater lattice energy.
79
How does ion size affect lattice energy?
Larger ion size (greater distance) results in smaller lattice energy.
80
What is the Born-Haber Cycle?
A thermodynamic cycle used to calculate lattice energy from the step-by-step analysis of energies involved in forming an ionic compound.
81
What rule do transition metals follow when forming ions?
They always lose electrons from the outermost s orbital first before losing from the inner d orbital.
82
What is oxidation in the context of transition metals?
The loss of electrons.
83
What is a covalent bond?
A bond formed when two non-metal atoms share one or more pairs of electrons.
84
What are the types of covalent bonds based on the number of shared electron pairs?
* Single Bond: One pair * Double Bond: Two pairs * Triple Bond: Three pairs
85
What is electronegativity?
A measure of an atom's ability to attract shared electrons within a chemical bond.
86
What is a Nonpolar Covalent Bond?
A bond where electrons are shared equally, typically between identical atoms.
87
What is a Polar Covalent Bond?
A bond where electrons are shared unequally, creating a dipole with partial charges.
88
What is the electronegativity difference that indicates an Ionic Bond?
ΔEN > 1.8.
89
How does molecular geometry affect molecular polarity?
Molecular geometry determines how individual bond dipoles add up, affecting overall polarity.
90
What characterizes a paramagnetic substance?
It has one or more unpaired electrons and is weakly attracted to a magnetic field.
91
What characterizes a diamagnetic substance?
All electrons are paired, and it is weakly repelled by a magnetic field.
92
What is ferromagnetism?
A strong form of paramagnetism found in materials like iron, where magnetic moments align, creating a strong magnet.
93
What is the significance of the octet rule?
Atoms tend to lose, gain, or share electrons to achieve 8 valence electrons.
94
What is a dipole moment?
A measure of the polarity of a bond or molecule.
95
What are the four main types of chemical bonds?
Ionic, covalent, metallic, and hydrogen bonds.
96
What is an ionic bond?
A bond formed between a metal and a non-metal through the complete transfer of electrons.
97
What happens to a metal atom in an ionic bond?
It loses its valence electrons and becomes a positively charged ion called a cation.
98
What is an example of a cation?
Sodium ion (Na⁺).
99
What happens to a non-metal atom in an ionic bond?
It gains electrons to become a negatively charged ion called an anion.
100
What is an example of an anion?
Chloride ion (Cl⁻).
101
What constitutes the ionic bond?
The strong electrostatic attraction between a cation and an anion.
102
What is lattice energy?
The energy released when gaseous ions form a stable crystal lattice.
103
What is a covalent bond?
A bond formed between two non-metal atoms through the sharing of electrons.
104
What characterizes a single covalent bond?
Two atoms share one pair of electrons.
105
What is an example of a single covalent bond?
Hydrogen molecule (H₂).
106
What characterizes a double covalent bond?
Two atoms share two pairs of electrons.
107
What is an example of a double covalent bond?
Oxygen molecule (O₂).
108
What characterizes a triple covalent bond?
Two atoms share three pairs of electrons.
109
What is an example of a triple covalent bond?
Nitrogen molecule (N₂).
110
What is electronegativity?
A measure of an atom's ability to attract shared electrons.
111
What is the most electronegative atom?
Fluorine (F).
112
What is a nonpolar covalent bond?
A bond with very small or zero difference in electronegativity between two atoms.
113
What is a polar covalent bond?
A bond with a significant difference in electronegativity, resulting in a partial charge.
114
What is the purpose of Lewis structures?
To represent the valence electrons of atoms and the covalent bonds within a molecule.
115
What is a bonding pair?
A pair of electrons shared between two atoms.
116
What is a lone pair?
A pair of valence electrons not involved in bonding.
117
What is an electron domain?
A region of electron density around a central atom.
118
What does VSEPR theory stand for?
Valence Shell Electron Pair Repulsion theory.
119
What does VSEPR theory state?
Electron domains repel each other and arrange themselves around the central atom.
120
What is an expanded octet?
A situation where atoms in Period 3 and below can accommodate more than eight electrons.
121
What is formal charge?
The hypothetical charge an atom would have if all bonding electrons were shared equally.
122
What is resonance in chemistry?
The occurrence of multiple valid Lewis structures differing only in electron placement.
123
What is a key feature of resonance structures?
The actual structure is a resonance hybrid of all possible structures.
124
In a covalent bond, what do the electrons do?
They are shared between atoms.
125
What is a resonance hybrid?
An average of multiple valid Lewis structures representing delocalized electrons
## Footnote
Resonance hybrids illustrate that electrons are not confined to a single bond or lone pair.
126
What is the total number of valence electrons in the nitrate ion (NO₃⁻)?
24 electrons
## Footnote
Calculated as: N = 5, O = 6, plus 1 for the -1 charge.
127
In the Lewis structure of nitrate ion, which atom is central?
Nitrogen (N)
## Footnote
Nitrogen is surrounded by three oxygen atoms.
128
How many equivalent resonance structures does the nitrate ion have?
Three equivalent resonance structures
## Footnote
The double bond can be formed with any of the three oxygen atoms.
129
What is the formal charge of the double-bonded oxygen in a resonance structure of nitrate?
0
## Footnote
The formal charges for the other atoms adjust to maintain the overall charge of -1.
130
How is bond order calculated in resonance structures?
Bond Order = (Total number of bonds between the two atoms in all structures) / (Total number of resonance structures)
## Footnote
This formula provides an average bond strength.
131
What is the bond order of the N-O bond in the nitrate ion?
1.33
## Footnote
Calculated from the bond types in the resonance structures.
132
What is the total number of valence electrons in ozone (O₃)?
18 electrons
## Footnote
Each oxygen contributes 6 valence electrons.
133
How many resonance structures does ozone have?
Two resonance structures
## Footnote
Each structure has a central oxygen bonded to two outer oxygens, with single and double bonds swapping positions.
134
What is the bond order for the O-O bond in ozone?
1.5
## Footnote
This indicates that the bonds are identical and intermediate between a single and double bond.
135
What does VSEPR Theory predict?
The 3D shape of molecules based on the number of electron domains
## Footnote
Shapes include linear, trigonal planar, tetrahedral, etc.
136
What is hybridization in chemistry?
The mixing of atomic orbitals to form hybrid orbitals
## Footnote
Explains molecular geometries, such as sp, sp², sp³ hybridization.
137
What does Molecular Orbital (MO) Theory describe?
Electrons belonging to the entire molecule rather than individual atoms
## Footnote
It helps explain phenomena like paramagnetism in O₂.
138
Define ionic bond.
A bond formed by the transfer of electrons between a metal and a non-metal
## Footnote
Metals form cations, while non-metals form anions.
139
Define covalent bond.
A bond formed by the sharing of electron pairs between non-metal atoms
## Footnote
These can be single, double, or triple bonds.
140
What is electronegativity?
The ability of an atom to attract shared electrons in a molecule
## Footnote
It determines the polarity of bonds.
141
Differentiate between polar and nonpolar bonds.
Polar bonds have unequal electron sharing; nonpolar bonds have equal sharing
## Footnote
This affects molecular properties and interactions.
142
What is a Lewis structure?
A diagram showing valence electrons as dots and bonds as dashes
## Footnote
It represents the molecular structure.
143
What is the octet rule?
Atoms tend to bond to achieve eight electrons in their valence shell
## Footnote
This rule guides many bonding scenarios.
144
What is an expanded octet?
An exception to the octet rule where elements can have more than 8 valence electrons
## Footnote
Common in elements in Period 3 and below.
145
What is a lone pair?
A pair of valence electrons not involved in bonding
## Footnote
These can affect molecular geometry.
146
What is an electron domain?
A region of electron density, including bonds or lone pairs
## Footnote
Determines molecular geometry.
147
What is formal charge?
A calculated charge on an atom in a molecule
## Footnote
Used to assess the validity of a Lewis structure.
148
What is resonance in chemistry?
Occurs when multiple valid Lewis structures can be drawn for a molecule
## Footnote
The true structure is a hybrid of these forms.
149
What does bond order indicate?
The average number of bonds between two atoms
## Footnote
It can often be a fractional value in resonance structures.
150
What are the four main types of chemical bonds?
Ionic, covalent, metallic, and hydrogen bonds.
151
What is an ionic bond?
A bond formed between a metal and a non-metal through the complete transfer of electrons.
152
What happens to a metal atom in an ionic bond?
It loses its valence electrons and becomes a positively charged ion called a cation.
153
What is an example of a cation?
Sodium ion (Na⁺).
154
What happens to a non-metal atom in an ionic bond?
It gains electrons to become a negatively charged ion called an anion.
155
What is an example of an anion?
Chloride ion (Cl⁻).
156
What constitutes the ionic bond?
The strong electrostatic attraction between a cation and an anion.
157
What is lattice energy?
The energy released when gaseous ions form a stable crystal lattice.
158
What is a covalent bond?
A bond formed between two non-metal atoms through the sharing of electrons.
159
What characterizes a single covalent bond?
Two atoms share one pair of electrons.
160
What is an example of a single covalent bond?
Hydrogen molecule (H₂).
161
What characterizes a double covalent bond?
Two atoms share two pairs of electrons.
162
What is an example of a double covalent bond?
Oxygen molecule (O₂).
163
What characterizes a triple covalent bond?
Two atoms share three pairs of electrons.
164
What is an example of a triple covalent bond?
Nitrogen molecule (N₂).
165
What is electronegativity?
A measure of an atom's ability to attract shared electrons.
166
What is the most electronegative atom?
Fluorine (F).
167
What is a nonpolar covalent bond?
A bond with very small or zero difference in electronegativity between two atoms.
168
What is a polar covalent bond?
A bond with a significant difference in electronegativity, resulting in a partial charge.
169
What is the purpose of Lewis structures?
To represent the valence electrons of atoms and the covalent bonds within a molecule.
170
What is a bonding pair?
A pair of electrons shared between two atoms.
171
What is a lone pair?
A pair of valence electrons not involved in bonding.
172
What is an electron domain?
A region of electron density around a central atom.
173
What does VSEPR theory stand for?
Valence Shell Electron Pair Repulsion theory.
174
What does VSEPR theory state?
Electron domains repel each other and arrange themselves around the central atom.
175
What is an expanded octet?
A situation where atoms in Period 3 and below can accommodate more than eight electrons.
176
What is formal charge?
The hypothetical charge an atom would have if all bonding electrons were shared equally.
177
What is resonance in chemistry?
The occurrence of multiple valid Lewis structures differing only in electron placement.
178
What is a key feature of resonance structures?
The actual structure is a resonance hybrid of all possible structures.
179
In a covalent bond, what do the electrons do?
They are shared between atoms.
180
What is a resonance hybrid?
An average of multiple valid Lewis structures representing delocalized electrons
## Footnote
Resonance hybrids illustrate that electrons are not confined to a single bond or lone pair.
181
What is the total number of valence electrons in the nitrate ion (NO₃⁻)?
24 electrons
## Footnote
Calculated as: N = 5, O = 6, plus 1 for the -1 charge.
182
In the Lewis structure of nitrate ion, which atom is central?
Nitrogen (N)
## Footnote
Nitrogen is surrounded by three oxygen atoms.
183
How many equivalent resonance structures does the nitrate ion have?
Three equivalent resonance structures
## Footnote
The double bond can be formed with any of the three oxygen atoms.
184
What is the formal charge of the double-bonded oxygen in a resonance structure of nitrate?
0
## Footnote
The formal charges for the other atoms adjust to maintain the overall charge of -1.
185
How is bond order calculated in resonance structures?
Bond Order = (Total number of bonds between the two atoms in all structures) / (Total number of resonance structures)
## Footnote
This formula provides an average bond strength.
186
What is the bond order of the N-O bond in the nitrate ion?
1.33
## Footnote
Calculated from the bond types in the resonance structures.
187
What is the total number of valence electrons in ozone (O₃)?
18 electrons
## Footnote
Each oxygen contributes 6 valence electrons.
188
How many resonance structures does ozone have?
Two resonance structures
## Footnote
Each structure has a central oxygen bonded to two outer oxygens, with single and double bonds swapping positions.
189
What is the bond order for the O-O bond in ozone?
1.5
## Footnote
This indicates that the bonds are identical and intermediate between a single and double bond.
190
What does VSEPR Theory predict?
The 3D shape of molecules based on the number of electron domains
## Footnote
Shapes include linear, trigonal planar, tetrahedral, etc.
191
What is hybridization in chemistry?
The mixing of atomic orbitals to form hybrid orbitals
## Footnote
Explains molecular geometries, such as sp, sp², sp³ hybridization.
192
What does Molecular Orbital (MO) Theory describe?
Electrons belonging to the entire molecule rather than individual atoms
## Footnote
It helps explain phenomena like paramagnetism in O₂.
193
Define ionic bond.
A bond formed by the transfer of electrons between a metal and a non-metal
## Footnote
Metals form cations, while non-metals form anions.
194
Define covalent bond.
A bond formed by the sharing of electron pairs between non-metal atoms
## Footnote
These can be single, double, or triple bonds.
195
What is electronegativity?
The ability of an atom to attract shared electrons in a molecule
## Footnote
It determines the polarity of bonds.
196
Differentiate between polar and nonpolar bonds.
Polar bonds have unequal electron sharing; nonpolar bonds have equal sharing
## Footnote
This affects molecular properties and interactions.
197
What is a Lewis structure?
A diagram showing valence electrons as dots and bonds as dashes
## Footnote
It represents the molecular structure.
198
What is the octet rule?
Atoms tend to bond to achieve eight electrons in their valence shell
## Footnote
This rule guides many bonding scenarios.
199
What is an expanded octet?
An exception to the octet rule where elements can have more than 8 valence electrons
## Footnote
Common in elements in Period 3 and below.
200
What is a lone pair?
A pair of valence electrons not involved in bonding
## Footnote
These can affect molecular geometry.
201
What is an electron domain?
A region of electron density, including bonds or lone pairs
## Footnote
Determines molecular geometry.
202
What is formal charge?
A calculated charge on an atom in a molecule
## Footnote
Used to assess the validity of a Lewis structure.
203
What is resonance in chemistry?
Occurs when multiple valid Lewis structures can be drawn for a molecule
## Footnote
The true structure is a hybrid of these forms.
204
What does bond order indicate?
The average number of bonds between two atoms
## Footnote
It can often be a fractional value in resonance structures.
205
What elements are in Period 2 of the periodic table?
Nitrogen (N), oxygen (O), and carbon (C)
## Footnote
These elements are characterized by their inability to exceed eight electrons in their valence shell.
206
What is the key rule regarding Period-2 elements and their bonding?
Period-2 elements cannot have more than 8 electrons in their valence shell and therefore cannot form more than four bonds.
207
What is the first principle for choosing a central atom in Lewis structures?
Carbon should be the central atom in many neutral covalent molecules because it typically forms four bonds and is less electronegative than N and O.
208
What should be done if octets are incomplete in a Lewis structure?
Introduce double bonds (e.g., C=N and C=O) to satisfy octets.
209
Can carbon form quadruple bonds with any single atom?
No, carbon cannot form a quadruple bond with any single atom in typical stable molecules.
210
How is formal charge calculated?
Formal charge = valence electrons − (nonbonding electrons + number of bonds).
211
What is the preferred structure criterion when computing formal charges?
Place negative formal charge on the most electronegative atom (generally O over N over C).
212
What are the three categories of octet-rule exceptions?
1. Species with an odd number of electrons (radicals)
2. Species with less than an octet (electron-deficient)
3. Species with more than eight electrons (hypervalent)
213
What is an example of an electron-deficient species?
BF3: Boron has six electrons and remains electron-deficient.
214
What is an example of a hypervalent species?
PCl5: Phosphorus in Period 3 can have 10 electrons around it.
215
Do fluorine and hydrogen form double bonds in standard neutral molecules?
No, fluorine and hydrogen rarely, effectively never, form double bonds.
216
Why does fluorine resist forming multiple bonds?
Because F is the most electronegative element, it strongly resists forming multiple bonds and keeps its lone pairs.
217
What is the recommended structure for phosphate-related species?
Prefer structures with single P–O bonds and negative charges on oxygens rather than forcing P–O double bonds.
218
How does bond order relate to bond length and strength?
As the number of bonds increases (single → double → triple), bond length decreases and bond strength increases.
219
What is the typical bond length for a C–C single bond?
≈ 1.54 Å
220
What is the typical bond length for a C=C double bond?
≈ 1.34 Å
221
What is the sign convention for bond enthalpy?
Breaking a bond requires energy input: positive sign (+); forming a bond releases energy: negative sign (−).
222
What is the formula for calculating reaction enthalpy via bond energies?
ΔH ≈ Σ(energies of bonds broken) − Σ(energies of bonds formed).
223
When counting bonds in reactions, what should be done for multiple identical bonds?
Multiply the bond enthalpy by the number of bonds breaking or forming.
224
What is resonance in the context of Lewis structures?
Resonance occurs when more than one valid Lewis structure can be drawn, and the real structure is a hybrid.
225
How can bond order be fractional in resonance hybrids?
Bond order can be fractional if a bond alternates between single and double across resonance structures.
226
What is the significance of bond enthalpy in reaction calculations?
Bond enthalpy is used to estimate the energy changes in chemical reactions.
227
What should be preferred when placing negative charges in oxyanions?
Prefer placing negative charge on oxygen in oxyanions.
228
What is the octet rule?
Atoms tend to have 8 electrons in their valence shell for stability.
229
What is hypervalency?
Period-3 and beyond can have expanded octets (e.g., 10 or 12 electrons) due to accessible d orbitals.
Chem 1411
flashcards
Decks in class (9)
# Cards
