Covalent Bonding & Structure (Edexcel International A Level (IAL) Chemistry): Flashcards

A dative covalent bond (also called a coordinate bond) is a covalent bond in which both electrons in the shared pair come from the same atom. The donor atom provides a lone pair to an electron-deficient acceptor atom.

One atom must have a lone pair of electrons available to donate, and the other atom must be electron-deficient, having an unfilled outer orbital that can accept the lone pair.

False.

In a dative covalent bond, both electrons come from the same atom (the donor). This distinguishes it from an ordinary covalent bond where each atom contributes one electron.

When NH₃ reacts with H⁺, nitrogen donates its lone pair to the electron-deficient proton. The H⁺ has an empty shell and accepts the pair, forming a dative covalent bond and giving the NH₄⁺ ion all four equivalent N-H bonds.

At lower temperatures, two molecules of AlCl₃ join together to form a dimer, Al₂Cl₆, using dative covalent bonds.

A single bond contains 1 shared pair of electrons, a double bond contains 2 shared pairs and a triple bond contains 3 shared pairs. More shared pairs result in a shorter and stronger bond.

False.

Covalent bonding occurs between non-metal atoms. Metals typically form ionic or metallic bonds rather than covalent bonds.

Only outer-shell (valence) electrons are shown. Dots represent electrons from one atom and crosses from the other. Shared pairs are drawn in the overlap between the two atoms, and lone pairs are shown on the outside of each atom.

An allotrope is one of two or more different physical forms of the same element, arising from different bonding arrangements of its atoms. For example, diamond and graphite are both allotropes of carbon.

In graphite, the layers of hexagonally arranged carbon atoms are held together only by weak van der Waals' forces. These forces are easily overcome, allowing the layers to slide over each other.

Each carbon atom in graphite bonds to only three others, leaving one delocalised electron per carbon atom. These delocalised electrons can move along the layers and carry charge when a voltage is applied.

False.

In diamond, all four outer electrons of every carbon atom are involved in covalent bonds. There are no delocalised electrons available to carry charge, so diamond does not conduct electricity.

In diamond, each carbon atom bonds covalently to four others in a tetrahedral arrangement, using all four outer electrons. In graphite, each carbon bonds to only three others in layers of hexagons, leaving one electron per atom delocalised.

Graphene is a single layer of graphite consisting of carbon atoms arranged in a continuous hexagonal layer, making it essentially a 2D molecule.

In diamond, each carbon atom is bonded to four others via strong covalent bonds in a rigid three-dimensional tetrahedral network. Breaking this structure requires overcoming a very large number of strong bonds simultaneously.

False.

Although both are allotropes of carbon, their different bonding arrangements give them completely different physical properties. Graphite is soft and conducts electricity, while diamond is hard and does not.

Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond towards itself. It is measured on the Pauling scale, with fluorine having the highest value of 4.0.

Fluorine has the highest electronegativity of any element, with a Pauling scale value of 4.0. It is the most effective at attracting bonding electrons towards itself.

False.

Although each C-Cl bond is polar, the four bonds are arranged symmetrically so the individual dipole moments cancel out. CCl₄ is therefore a nonpolar molecule overall.

A difference of 2 or more on the Pauling scale leads to an ionic bond. A difference of less than 2 leads to a covalent bond, with greater polarity as the difference increases toward 2.

Dipole moment is a measure of how polar a bond is. Its direction is shown by an arrow pointing toward the partially negatively charged (δ−) end of the bond.

A molecule is overall polar if the individual dipole moments of its polar bonds do not cancel each other out. This depends on both the polarity of each bond and the shape and symmetry of the molecule.

In a polar covalent bond, the more electronegative atom gains a partial negative charge, written as δ−.

A nonpolar covalent bond exists when two atoms with the same electronegativity share electrons equally. The electron distribution is symmetric and there are no partial charges on either atom.

In CCl₄ the four C-Cl dipoles are arranged symmetrically and cancel out, giving a net dipole of zero. In CH₃Cl the four different bonds are arranged asymmetrically, so the dipoles do not cancel and the molecule has an overall dipole moment.

False.

A difference of 2 or more on the Pauling scale leads to an ionic bond, not a covalent one. Covalent bonds form when the electronegativity difference is less than 2.

VSEPR theory (valence shell electron pair repulsion theory) predicts the shape and bond angles of molecules. It states that electron pairs in the valence shell of the central atom repel each other and arrange themselves to minimise these repulsions.

As bond order increases, bond length decreases and bond strength increases. Greater electron density between the nuclei increases the attractive forces, pulling the atoms closer together and making the bond harder to break.

False.

Lone pair–lone pair repulsion is the strongest type. Lone pairs have a more concentrated electron charge cloud that sits closer to the central nucleus, exerting greater repulsion than bonding pairs.

The order of repulsion is: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair. Lone pairs exert the greatest repulsion because their electron charge cloud is more concentrated and closer to the central atom.

Bond length is the internuclear distance between two covalently bonded atoms. Shorter bonds arise from greater forces of attraction between electrons and nuclei, and shorter bonds are also stronger bonds.

Lone pairs are not shared between two nuclei, so their electron charge cloud is more concentrated and sits closer to the central atom's nucleus. This gives lone pairs a stronger repulsive effect on neighbouring electron pairs.

Triple bonds are the shortest and strongest covalent bonds due to the high electron density between the nuclei.

A molecule adopts the shape that minimises the repulsion between electron pairs in the valence shell of the central atom. This is the most stable arrangement for the bonding and lone pair electrons.

False.

According to VSEPR theory, a molecule adopts the shape that minimises repulsion between electron pairs. This gives the most stable arrangement of electrons around the central atom.

BeCl₂ is linear with a bond angle of 180°. Beryllium has 2 valence electrons, both used in bonding with no lone pairs, so the two bond pairs arrange linearly to minimise repulsion.

PCl₅ in the gas phase has a trigonal bipyramidal shape with bond angles of 120° and 90°.

BCl₃ is trigonal planar with bond angles of 120°. Boron has 3 valence electrons, all used in bonding with no lone pairs, so the three bond pairs spread evenly in a plane.

True.

Both NH₄⁺ and CH₄ are tetrahedral with bond angles of 109.5°. In NH₄⁺, all four N-H bonds (including the dative bond) are equivalent and there are no lone pairs to distort the angle.

CH₄ is tetrahedral with bond angles of 109.5°. Carbon has 4 valence electrons, all used in bonding with no lone pairs, so the four bond pairs arrange in a tetrahedral geometry.

NH₃ is a trigonal pyramid with bond angles of 107°. Nitrogen has 3 bonding pairs and 1 lone pair. The lone pair exerts greater repulsion than a bonding pair, compressing the H-N-H angles below the 109.5° of CH₄.

False.

BCl₃ is trigonal planar with bond angles of 120°. Boron has 3 bonding pairs and no lone pairs, so the three pairs spread evenly at 120° rather than the 109.5° of a tetrahedral arrangement.

NH₄⁺ is tetrahedral with bond angles of 109.5°. All four bonds (three N-H covalent bonds and one dative bond from the lone pair) are equivalent, giving the same geometry as CH₄.

Each carbon atom in C₂H₄ is trigonal planar. Each carbon has three regions of electron density (two C-H bonds and one C=C double bond), which arrange at 120° to minimise repulsion.

True.

SF₆ has 6 bonding pairs and no lone pairs. The six bond pairs arrange in an octahedral shape with all bond angles equal to 90°.

H₂O is bent (V-shaped) with a bond angle of 104.5°. Oxygen has 2 bonding pairs and 2 lone pairs. Lone pairs repel more strongly than bonding pairs, compressing the bond angle below 109.5°.

CO₂ is linear with a bond angle of 180°. It has two C=O double bonds and no lone pairs on carbon, so the two regions of electron density arrange linearly to minimise repulsion.