7.06 Quiz: Hydrogen & Metallic Bonding Practice

Hydrogen bonds are intermolecular forces that occur when a hydrogen atom attached to a highly electronegative element interacts with another electronegative atom. They are much weaker than covalent bonds and influence properties like boiling and melting points.

Water is a polar molecule with O-H bonds that allow for hydrogen bonding between molecules. This network of hydrogen bonds is largely responsible for water's high boiling point and unique physical properties.

Metallic bonding involves a structure where metal atoms are arranged in a lattice and the outer electrons are free to move throughout. This sea of delocalized electrons gives metals their hallmark properties such as electrical conductivity and malleability.

The presence of delocalized electrons in metallic bonds allows electrons to move freely, which is why metals conduct electricity so well. This same electron mobility also contributes to other properties like thermal conductivity and malleability.

Hydrogen bonding occurs when hydrogen is bound to highly electronegative elements that pull electron density towards themselves. Fluorine, oxygen, and nitrogen are common examples of such elements, making them ideal partners for forming hydrogen bonds.

Hydrogen bonds create additional attractions between water molecules, requiring more energy to separate them during boiling. This extra energy requirement results in water having a higher boiling point compared to similar-sized molecules without hydrogen bonding.

In metallic bonding, the electrons are not tied to any one atom which allows the atoms to shift position relative to each other. This freedom of movement results in metals being malleable and ductile.

Hydrogen bond strength is primarily determined by factors such as electronegativity, bond length, and bond angle. The small, nearly constant size of the hydrogen atom does not have a significant influence on bond strength.

Metallic bonds consist of a lattice of metal ions immersed in a sea of delocalized electrons. This structure is key to many of the unique properties of metals, such as electrical conductivity.

Strong metallic bonds arise from the attraction between metal ions and the delocalized sea of electrons. Breaking these bonds requires significant energy, which contributes to the high melting points of metals.

Hydrogen bonds are intermolecular forces that occur between distinct molecules, whereas metallic bonds are found within a continuum of metal atoms. This fundamental difference explains why metals exhibit properties like conductivity and malleability.

More hydrogen bonds mean stronger intermolecular attractions, which require additional energy to overcome when changing phases. Consequently, substances with extensive hydrogen bonding typically exhibit higher boiling points and increased viscosity.

The delocalized electrons in metallic bonds move freely through the metal lattice, allowing efficient transfer of heat and electricity. This same phenomenon also permits atoms to slide past one another, lending metals their malleability.

The free movement of delocalized electrons in metals is responsible for their excellent thermal and electrical conductivity. Additionally, this electron mobility allows the atoms to rearrange without breaking bonds, making metals both strong and malleable.

Hydrogen bonding arises due to the unequal distribution of electrons in a polar bond, creating a partial positive charge on the hydrogen and a partial negative charge on the electronegative atom. This electrostatic attraction is a specialized case of dipole-dipole interactions.

Covalent bonds are strong and require substantial energy to break, whereas hydrogen bonds are significantly weaker. This difference allows hydrogen bonds to form and break easily, playing a key role in dynamic processes such as phase changes and biological interactions.

In alloys, the mixture of different metal atoms in a common lattice allows for a shared pool of delocalized electrons. This electron mobility can enhance conductivity, strength, and other physical properties compared to pure metals.

Hydrogen bonds are critical for maintaining the alpha-helices and beta-sheets in protein secondary structures. Disruption in these bonds can lead to misfolding or denaturation, ultimately affecting the protein's biological function.

As temperature rises, atoms in a metal vibrate more intensely, which in turn disrupts the orderly movement of delocalized electrons. This increased lattice vibration can scatter electrons, reducing overall metallic bond strength and affecting the metal's conductivity.

The non-directional nature of metallic bonds lets metal atoms slide past each other without bond breakage due to the delocalized electron cloud. In contrast, ionic compounds have a rigid lattice where displacement of ions can bring like charges into close proximity, causing repulsive forces and brittleness.