Covalent Bonding Concepts

False (B)

Sharing electrons

<p>cross</p> Signup and view all the answers

<p>H3O+ = Hydronium ion NH4+ = Ammonium ion Cl2 = Chlorine molecule H2 = Hydrogen molecule</p> Signup and view all the answers

<p>Energy is released during bond formation. (C)</p> Signup and view all the answers

<p>True (A)</p> Signup and view all the answers

<p>They determine the chemical behavior of the element.</p> Signup and view all the answers

<p>They are not associated with a single atom or one covalent bond. (A)</p> Signup and view all the answers

<p>False (B)</p> Signup and view all the answers

<p>resonance hybrid</p> Signup and view all the answers

<p>expanded octet</p> Signup and view all the answers

<p>Incomplete Octet = Too few valence electrons present Expanded Octet = Too many valence electrons present Odd Number of Electrons = Uneven distribution of valence electrons</p> Signup and view all the answers

<p>Has bonds that are all of equal length (B), Has one double bond and two single bonds (D)</p> Signup and view all the answers

<p>False (B)</p> Signup and view all the answers

<p>formal charges</p> Signup and view all the answers

<p>Bent or V-shaped (D)</p> Signup and view all the answers

<p>False (B)</p> Signup and view all the answers

<p>Sigma bond</p> Signup and view all the answers

<p>tetrahedral</p> Signup and view all the answers

<p>BeH2 = Linear NH3 = Trigonal pyramidal SO2 = Bent PCl5 = Trigonal bipyramidal</p> Signup and view all the answers

<p>SO2 (B)</p> Signup and view all the answers

<p>False (B)</p> Signup and view all the answers

<p>Trigonal planar</p> Signup and view all the answers

<p>Bond strength increases (D)</p> Signup and view all the answers

<p>False (B)</p> Signup and view all the answers

<p>di</p> Signup and view all the answers

<p>shorter</p> Signup and view all the answers

<p>BF3 (D)</p> Signup and view all the answers

<p>mono = 1 di = 2 tri = 3 tetra = 4</p> Signup and view all the answers

<p>an arrow</p> Signup and view all the answers

<p>False (B)</p> Signup and view all the answers

<p>Tetrahedral (A)</p> Signup and view all the answers

<p>False (B)</p> Signup and view all the answers

<p>109.5</p> Signup and view all the answers

<p>linear</p> Signup and view all the answers

<p>2 = Linear 3 = Trigonal Planar 4 = Tetrahedral 5 = Trigonal Bipyramidal 6 = Octahedral</p> Signup and view all the answers

<p>Trigonal Bipyramidal (D)</p> Signup and view all the answers

<p>False (B)</p> Signup and view all the answers

<p>Trigonal Pyramidal</p> Signup and view all the answers

<p>octahedral</p> Signup and view all the answers

<p>Linear = 2 electron domains Trigonal Planar = 3 electron domains Tetrahedral = 4 electron domains Octahedral = 6 electron domains</p> Signup and view all the answers

<p>The sharing of electrons between atoms (A)</p> Signup and view all the answers

<p>False (B)</p> Signup and view all the answers

<p>The measure of the force of attraction of an atom towards the electrons in a bond.</p> Signup and view all the answers

<p>bent</p> Signup and view all the answers

<p>Single bond = Two electrons shared Double bond = Four electrons shared Triple bond = Six electrons shared Dative bond = One atom donates a pair of electrons</p> Signup and view all the answers

<p>BF3 (D)</p> Signup and view all the answers

<p>False (B)</p> Signup and view all the answers

<p>A bond where electrons are shared unequally between two atoms, leading to a partial charge.</p> Signup and view all the answers

The strength of a covalent bond is directly proportional to the number of electron pairs shared between two atoms. More shared electrons result in a stronger pull between the atoms, causing the bond to be shorter.

A covalent bond where both electrons in the shared pair are contributed by one atom. This results in a positive charge on one atom and negative on the other.

The naming of molecules uses a prefix system to indicate the number of atoms present within each molecule.

The difference in electronegativity between two bonded atoms determines the polarity of the bond.

A scale that ranges from ionic to covalent bonding. Although ionic and covalent bonds are considered distinct, they are actually extremes on a spectrum.

A type of chemical bond where two non-metal atoms share electrons to achieve a stable outer shell configuration.

The outermost shell of an atom, which contains the electrons involved in chemical bonding.

Symbols that represent atoms and only show their valence electrons.

The rule that most atoms strive to have eight electrons in their valence shell to achieve stability.

The force of attraction between the nuclei of two atoms and the shared electrons in a covalent bond.

The distance between the nuclei of two atoms in a covalent bond, representing the optimal balance of attractive and repulsive forces.

The energy released during the formation of a covalent bond, indicating its stability.

The process of showing the arrangement of electrons in a molecule, using dots or lines to represent electron pairs.

A covalent bond is formed when two atoms share one or more electrons. This sharing occurs between atoms with similar electronegativity values and results in an electrostatic attraction between the positively charged nuclei and the negatively charged shared electrons.

A dative or co-ordinate covalent bond forms when one atom provides both electrons for a shared pair. This occurs when one atom has a lone pair of electrons and the other atom has an empty orbital.

Electronegativity is the measure of an atom's ability to attract electrons in a chemical bond. It is a relative value, with higher electronegativity indicating a stronger pull towards electrons.

When the electronegativity difference between two atoms in a bond is small or zero, the bond is considered non-polar. The electrons are shared equally between the atoms.

A polar covalent bond occurs when the electronegativity difference between two atoms is significant, resulting in an uneven sharing of electrons. This creates a partial positive charge (δ+) on the less electronegative atom and a partial negative charge (δ-) on the more electronegative atom.

The VSEPR (Valence Shell Electron Pair Repulsion) theory predicts the shape of molecules based on the repulsion between electron pairs in the valence shell of the central atom. This repulsion causes electron pairs to arrange themselves as far apart as possible, minimizing their interactions.

Valence bond theory describes covalent bonding in terms of the overlap of atomic orbitals. When two atomic orbitals overlap, they form a new molecular orbital that can hold two electrons. This overlap can be head-to-head (sigma bond) or side-to-side (pi bond).

Electrons that aren't tied to a specific atom or bond in a molecule, ion, or metal. They move freely within the structure.

Different Lewis structures that can be drawn for the same molecule or ion, each showing a different distribution of electrons.

The true structure of a molecule or ion, an average of all its possible resonance structures. It shows delocalized electrons and bond lengths that are intermediate between single and double bonds.

The structure of a molecule or ion that minimizes formal charges and places negative charges on the most electronegative atoms.

When a molecule's central atom has less than 8 electrons in its valence shell.

When a molecule's central atom has more than 8 electrons in its valence shell.

When a molecule has an odd number of valence electrons, making it impossible for all atoms to have a full octet.

The total number of electron groups (bonds and lone pairs) surrounding the central atom in a molecule or ion.

The arrangement of electron domains around the central atom, regardless of the type of bond.

The actual shape of a molecule or ion, taking into account the repulsions between bonding pairs and lone pairs.

Two electron groups around the central atom result in a linear arrangement.

Three electron groups around the central atom result in a flat, triangular arrangement.

Four electron groups around the central atom result in a pyramid-like arrangement.

Five electron groups around the central atom result in a shape with two pyramids connected at their bases.

Six electron groups around the central atom result in a symmetrical, cube-like arrangement.

Lone pairs exert greater repulsion than bonding pairs, causing a distortion in the molecular shape.

The shape of a molecule depends on the number of lone pairs and bonding pairs around the central atom.

The Valence Shell Electron Pair Repulsion (VSEPR) theory provides the shape of molecules and ions, but offers no information regarding the orbitals within the molecule or the characteristics of covalent bonds. The Valence Bond theory focuses on answering these specific questions.

Sigma bonds are characterized by symmetrical electron density along the axis connecting the nuclei of the bonded atoms, resulting in a strong electrostatic attraction that holds the atoms together.

A pi bond is a type of covalent bond where the electron density lies above and below the plane of the bonded atoms. Unlike sigma bonds, electrons are delocalized in a pi bond, leading to weaker electrostatic attraction between the electrons and the nuclei. This also makes pi bonds more susceptible to breaking.

The spatial arrangement of the lone pairs and bonded atoms surrounding a central atom is called the molecular geometry.

Hybrid orbitals are formed by mixing the atomic orbitals of the same atom.

Sp3 hybridization involves the mixing of one s and three p orbitals to produce four equivalent sp3 hybrid orbitals. This hybridization occurs for central atoms that have four electron domains (four bonds or two bonds and two lone pairs). sp3 hybrid orbitals have a tetrahedral shape, resulting in a tetrahedral geometry of the molecule. The presence of lone pairs can alter the shape of the molecule from true tetrahedral geometry to trigonal pyramidal or bent.

Hybrid orbitals are formed by mixing the atomic orbitals of the same atom.

The geometry of sp3 hybridization is tetrahedral. This means that the four hybrid orbitals point towards the corners of a tetrahedron, resulting in a tetrahedral molecular geometry when the central atom only forms bonds. However, the presence of lone pairs on the central atom can distort the shape of the molecule.

For instance, if the central atom has three bonds and one lone pair, the molecular geometry will be trigonal pyramidal. If the central atom has two bonds and two lone pairs, the molecular geometry will be bent. This distortion happens because the lone pairs exert more repulsion than bonding pairs, pushing the bonded atoms closer to each other.