Organic vs Inorganic Chemistry Quiz

Organic chemistry is the branch of chemistry that focuses on the study of carbon-containing compounds. It involves understanding the structure, properties, reactions, and synthesis of these compounds.

The historical belief that organic compounds could only be made by living organisms. This theory was debunked by scientists like Chevreul and Woehler.

In 1828, Friedrich Wöhler demonstrated that urea, a naturally occurring organic compound found in urine, could be synthesized from inorganic ammonium cyanate. This groundbreaking experiment challenged the vital force theory and paved the way for synthetic organic chemistry.

Carbon can form four covalent bonds due to having four valence electrons. This ability to form multiple bonds with other atoms allows for the creation of a vast array of complex and diverse molecules.

Organic compounds are far more numerous than inorganic compounds because carbon's ability to form diverse structures with other atoms is highly efficient, leading to an extensive variety of molecular possibilities.

Carbon atoms can form more than one unique compound with the same molecular formula but different structural arrangements.

Covalent bonds between carbon and hydrogen are strong and stable.

Carbon has four valence electrons, allowing it to form four covalent bonds.

Kekule and Couper observed that carbon always forms four bonds. Later, van't Hoff and Le Bel proposed that these bonds have specific spatial directions.

Atoms bond to form a more stable compound than individual atoms. This stability is achieved through ionic or covalent bonds, depending on the atom type.

The sharing of electrons between atoms to form a stable molecule with a complete outer shell of electrons. In most atoms, this means having 8 valence electrons (octet rule).

Visual representations of electron pairs and single bonds, used to depict the valence electron arrangement in a molecule.

Line-bond structures that represent covalent bonds with lines and skeletal formulas without explicitly drawing hydrogen atoms.

Atoms with four or more valence electrons tend to form as many bonds as needed to complete their valence shell to achieve a stable octet (eight electrons).

Carbon has four valence electrons and forms four bonds to complete its octet, as seen in methane (CH4).

Nitrogen has five valence electrons but only forms three bonds in ammonia (NH3), leaving one lone pair of electrons.

Oxygen has six valence electrons and forms two bonds in water (H2O) with two lone pairs of electrons remaining.

Valence electrons that are not involved in bonding are called nonbonding electrons, or lone-pair electrons.

A covalent bond forms when two atoms share electrons to achieve a stable configuration. It results from the overlap of singly occupied orbitals from each atom.

Valence bond theory describes covalent bonding by focusing on the overlap of atomic orbitals to form shared electron pairs between atoms.

The energy required to break a given bond is called bond energy. A stronger bond requires more energy to break.

The number of atoms of each element in one molecule of a compound, providing no information about the arrangement of atoms.

The simplest whole-number ratio of elements in a compound. It represents the relative proportions, not the actual number of atoms.

A structural representation showing all atoms and bonds in a molecule, with each bond explicitly drawn.

A structural representation showing all atoms in a molecule, but only showing bonds when necessary to avoid redundancy.

A simplified structure showing only the carbon-carbon bonds and non-hydrogen atoms. Hydrogen atoms are implied.

A type of covalent bond formed by the direct overlap of atomic orbitals. It has a circular cross-section and is typically strong.

A type of covalent bond formed by the sideways overlap of p orbitals. It has a "dumbbell" shape and is typically weaker than a sigma bond.

A type of covalent bond that involves the overlap of atomic orbitals to create a new hybrid orbital. This process results in a specific geometry for the molecule.

A type of hybridization where one s orbital and three p orbitals combine to form four equivalent sp³ orbitals, resulting in a tetrahedral shape.

The simplest alkane with the chemical formula CH4, where carbon has four sp³ orbitals that overlap with the 1s orbitals of four hydrogen atoms, resulting in four identical C-H bonds.

The molecule ethane is formed by the overlap of two sp³ orbitals from each carbon atom to form a C-C bond. The remaining sp³ orbitals on each carbon atom overlap with the 1s orbitals of hydrogen atoms to form six C-H bonds.

One s orbital and two p orbitals combine to form three sp² orbitals, which lie in a plane with 120° angles, while the remaining p orbital is perpendicular to the plane.

A carbon atom can form a double bond with another carbon atom. One bond (s bond) is formed by overlapping sp² orbitals, and the second bond (π bond) is formed by sideways overlap between two p orbitals.

The molecule ethylene (C2H4) has two C atoms connected by a C=C double bond. Each C atom is sp² hybridized and forms two C-H sigma bonds with two hydrogen atoms.

A carbon atom forms a triple bond with another carbon atom, where one bond (sigma bond) involves overlap of sp orbitals, and the other two bonds (π bonds) are formed by sideways overlap of two p orbitals.

This molecule (C₂H₂) has a linear structure with two C atoms connected by a triple bond. The carbon atom in each molecule is sp hybridized. The two sp orbitals form a σ bond between the C atoms.

A type of hybrid orbital created by combining one s and one p orbital, resulting in two sp hybrid orbitals that are linear and 180° apart.

Nitrogen and Oxygen can also form hybrid orbitals to form bonds.

A description of the probability of finding an electron at a specific location in a molecule, considering the energy levels and shapes of the orbitals.

Describes the probability of finding an electron in a specific region of space in a molecule, which results from the combination of atomic orbitals.

A molecular orbital that results from the constructive interference of atomic orbitals, which leads to a lower energy level and increased electron density between the atoms.

A molecular orbital that results from the destructive interference of atomic orbitals, leading to a higher energy level and decreased electron density between the atoms.