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| Understanding the concepts
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| 1.a.
| An ionic bond is the attraction of two oppositely charged ions for each other. Electrons are transferred from one atom to another when an ionic bond forms between a metal and a nonmetal.
Covalent bonding is the sharing of electrons between two atoms. Covalent bonding results when a bond forms between two nonmetals. Polar covalent bonding is when there is unequal sharing of the bonding electrons between two atoms. Polar covalent bonds usually form when two different nonmetals form a bond. The most electronegative atom in the bond is the partial negative end of the bond dipole, while the least electronegative atom is the partial positive end.
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| b.
| The bond length distance in the plot is the point where energy is at its lowest value. As atoms are brought closer together than the bond length distance, energy increases because the two positively charged nuclei of each atom start interacting with each other.
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| 2.a.
| Electronegativity is the ability of an atom in a molecule to attract shared electrons to itself. Electronegativity values increase
left to right across the periodic table
and decrease top to bottom.
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| b.
| Hydrogen has an electronegativity value between boron and carbon and has a value identical to phosphorus and tellurium. |
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| 3.a.
| Metals form stable cations and nonmetals form stable anions. The charged ions that various representative elements form can be predicted from the group location. The alkali metals all form stable +1 ions, the alkaline earth metals all form stable +2 ions, and metals in Group 3A form stable +3 ions. The nonmetals in the halogen group form stable -1 ions, the nonmetals in the oxygen group form stable -2 ions, and the nonmetals in the nitrogen group form stable -3 ions. The reason for the specific charged ions is the electron configuration. By forming these charged ions, the atoms in the various groups achieve a stable noble gas electron configuration.
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| b.
| Cations are smaller than the parent atom; anions are larger that the parent atom. |
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| 4.b.
| For CO2, each oxygen should only have 2 lone pairs (not 3). For NO2-, O3, and SO2, the terminal oxygen with the single bond should have 3 lone pairs (not 2), and the terminal oxygen with the double bond should have 2 lone pairs (not 3).
Resonance occurs when more than one valid Lewis structure can be drawn for a molecule or ion. Resonance structures usually involve different placements of multiple bonds. NO2-, O3, and SO2 all have two resonance structures. In each case, the central atom can satisfy the octet rule by forming a double bond to one oxygen and a single bond to the other oxygen. CO32- and SO3 each have three resonance structures (the central atom forms a double bond to one of the oxygens and the other two bonds to oxygen are single bonds).
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| c.
| BH3 has only six valence electrons and BeH2 has only four valence electrons, which makes it impossible to satisfy the octet rule for B and Be in these compounds. In general, B and Be Lewis structures often have fewer than eight electrons around them. For NO and NO2, there are an odd number of valence electrons, which also makes it impossible to satisfy the octet rule for all atoms. |
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| 5.a.
| The main idea of the VSEPR model is that the structure around a given atom is determined by minimizing repulsions between electron pairs. The geometry that places two pairs of electrons as far apart as possible is a linear arrangement with a 180° bond angle; three pairs of electrons occupy a trigonal planar arrangement with 120° bond angles; four pairs of electrons occupy a tetrahedral arrangement with 109.5° bond angles. |
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| b.
| When lone pairs of electrons are present around an atom, the molecular structure is not the same as the electron pair arrangement, e.g., O3 has a bent molecular structure (not a trigonal planar structure), NH3 has a trigonal pyramid structure (not a tetrahedral structure), and H2O has a bent molecular structure (not a tetrahedral structure). |
