Ch.8 - Basic Concepts of Chemical Bonding

Brown - Chemistry: The Central Science 14th Edition

Brown14th EditionChemistry: The Central ScienceISBN: 9780134414232당신이 사용하는 게 아니라요?교과서 변경

Brown 14th Edition

Ch.8 - Basic Concepts of Chemical Bonding

문제 48

8장, 문제 48

Write Lewis structures for the following: (a) H2CO (both H atoms are bonded to C), (b) H2O2, (c) C2F6 (contains a C¬C bond), (d) AsO33 - , (e) H2SO3 (H is bonded to O), (f) NH2Cl.

검증된 단계별 안내

Step 1: For H2CO, start by placing the carbon atom in the center as it is the least electronegative. Connect each hydrogen atom to the carbon. Then, place the oxygen atom connected to the carbon. Complete the octet for carbon and oxygen by adding double bonds or lone pairs as needed.

Step 2: For H2O2, place the two oxygen atoms in the center and connect them with a single bond. Attach each hydrogen to an oxygen atom. Ensure each oxygen has a total of eight electrons by adding lone pairs.

Step 3: For C2F6, start by placing the two carbon atoms in the center connected by a single bond. Attach three fluorine atoms to each carbon atom. Ensure each carbon atom has four bonds.

Step 4: For AsO33-, place the arsenic atom in the center and connect it to three oxygen atoms. Since it is an ion, ensure the total charge is -3 by adding lone pairs to the oxygen atoms.

Step 5: For H2SO3, place the sulfur atom in the center, connect it to two oxygen atoms with double bonds and one oxygen atom with a single bond. Attach the hydrogen atom to the singly bonded oxygen. Complete the structure by adding lone pairs to achieve stable octets and formal charges.

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Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms in a molecule and the lone pairs of electrons that may exist. They use dots to represent valence electrons and lines to represent covalent bonds. Understanding how to draw Lewis structures is essential for visualizing molecular geometry and predicting the behavior of molecules in chemical reactions.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in determining how atoms bond with each other. The number of valence electrons influences the molecule's reactivity and stability. In Lewis structures, it is important to account for all valence electrons to ensure that the structure accurately represents the molecule's electron configuration.

Formal Charge

Formal charge is a concept used to determine the distribution of electrons in a molecule and to assess the stability of a Lewis structure. It is calculated by comparing the number of valence electrons in the free atom to the number of electrons assigned to it in the Lewis structure. A lower formal charge on atoms generally indicates a more stable structure, guiding chemists in choosing the most favorable Lewis representation.

관련 실천

교과서 질문

Draw Lewis structures for the following: (f) NH₂OH

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교과서 질문

Draw Lewis structures for the following: (d) H₂SO₄ (H is bonded to O) (e) OF₂

Which one of these statements about formal charge is true? (a) Formal charge is the same as oxidation number. (b) To draw the best Lewis structure, you should minimize formal charge. (c) Formal charge takes into account the different electronegativities of the atoms in a molecule. (d) Formal charge is most useful for ionic compounds. (e) Formal charge is used in calculating the dipole moment of a diatomic molecule.

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교과서 질문

Draw Lewis structures for the following: (a) CH₂Cl₂ (b) ClCN (c) SF₂

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교과서 질문

For each of the following molecules or ions of sulfur and oxygen, write a single Lewis structure that obeys the octet rule, and calculate the oxidation numbers and formal charges on all the atoms: (a) SO₂ Write a single Lewis structure that obeys the octet rule for SO₂ and assign the formal charges on all the atoms. (b) SO₃ Write a single Lewis structure that obeys the octet rule for SO₃ and assign the formal charges on all the atoms. (c) SO₃^2- Write a single Lewis structure that obeys the octet rule for SO₃^2- and assign the formal charges on all the atoms.

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