HS Chemistry - Chemical Bonding

Unit Summary

Ionic Bonding

• Ionic bonds are formed between metals and non-metals. The metal becomes a cation, giving up its valence electrons until it achieves a noble gas configuration. The non-metal becomes an anion, taking electrons until it achieves a noble gas configuration. When both achieve a noble gas configuration $not necessarily with a 1:1 ratio of metal atoms to non-metal atoms$, an ionic compound can be formed $NaCl, AlF~3~, Li~2~O, etc.$.

• Metals have low 1^st ionization energies, so they can easily lose valence electrons. Non-metals have high electronegativities, so they can easily pull the electrons away from the metals.

• In an ionic compound, each ion is surrounded by oppositely charged ions, which it both attracts and is attracted to. This structure is called an ionic crystal. Thus, there is no distinct ionic molecule.

• In ionic compounds, the formula tells us the ratio of ions, not the total number.

• In an ionic bond, every ion is surrounded by an oppositely charged ion, which it both attracts and is attracted to.

• Ionic compounds have high melting and boiling points because the bonds between any two oppositely charged ions is strong, and there are multiple bonds holding each ion in place. Therefore, more force is required to break these bonds.

• This structure also makes ionic compounds incredibly brittle.

• Normally, the ions in an ionic compound are fixed in their positions. Since they can't move from one end of the substance to another, they can't carry a charge from one end of the substance to another, and therefore ionic compounds normally can't conduct electricity. However, when the ionic compound is dissolved or in the liquid state, the ions can move from one end of the substance to another, and therefore ionic compounds can conduct electricity in those states.

Covalent Bonding

• Covalent bonds usually form between two non-metals. In a covalent bond, the atoms share electrons to try and get a full octet. The shared pair of electrons is called the bonding pair. The atoms are really trying to pull the electrons towards themselves in order to complete their outer shell, but the nuclei of the bonded atoms are attracted by the electrons, so they don't leave each other. They also repel each other, so they don't get too close. This results in a fixed molecule shape. Covalent bonds form distinct molecules.

• Coordinate covalent bonds are covalent bonds where the entire electron pair is donated by one atom. Although they otherwise function the same as normal covalent bonds, they only form when the atoms/molecules bonding together have opposite dipoles.

• There are a few violations of the octet rule:

  • Odd-electron molecules: Stable molecules in which there are an odd number of electrons in the valence shell, showing that at least one atom doesn't have 8 valence electrons. An example is NO.
  • Electron-deficient molecules: Stable molecules in which one atom has less than 8 valence electrons. An example is BeCl₂.
  • Expanded Valence Shell Molecules: Stable molecules in which one atom has more than 8 electrons in its valence shell. An example is SF₆.

Metallic Bonding

• Metallic bonds are formed between metal atoms. In metallic bonding, all the valence electrons are delocalized, and travel freely across all the atoms connected by the metallic bonds. This happens because they are very tightly packed together, so their valence shells intersect those of multiple other atoms, and because they have less than 4 valence electrons each, so they have space for more valence electrons. These two things are unique to metallic bonding. There are no distinct molecules in metallic bonding.

• Some properties of metals:
  • Good conductors of electricity
  • Metallic bonds don't break easily.
  • High melting and boiling points.
  • Good conductors of heat.
  • Very malleable.
  • Lustrous $shiny$
  • Insoluble in water

• An alloy is a mixture of chemical elements in which at least one is a metal. They retain most of the characteristics of metals, but aren't malleable, as they contain atoms of different sizes. These disrupt the metal's regular structure, preventing the atoms from easily sliding over other atoms.

• When Transition Metals form metallic bonds, the d-subshell electrons also become delocalized, which changes some of their characteristics.

Allotropes

• Allotropes are different molecular arrangements of the same element.

• Allotropes have different molecular structures.

• Carbon has three main allotropes:

  • In graphite, the most stable form of carbon, the carbon atoms are arranged in layers of hexagonal patterns. Each carbon atom is bonded to three others.

    • The multiple bonds each carbon atom has gives graphite a high melting and boiling point.
    • Since graphite is arranged in layers that are connected by weak van der Waals forces, the forces can easily be broken, and thus graphite can easily be cut.
    • Since each carbon atom is bonded to three others, each carbon atom has one delocalized electron, which allows graphite to conduct electricity.

  • In diamond, the carbon atoms are arranged in a crystalline structure, with each carbon atom bonded to four others.

    • The multiple bonds each carbon atom has gives diamond a high melting and boiling point.
    • Since diamond has a crystalline structure, and each carbon atom is bonded to four others, diamond is very hard.

  • In fullerenes, the carbon atoms are arranged in hexagonal patterns. However, unlike graphite, there aren't multiple layers of hexagons.

    • Since fullerenes don't have multiple layers of hexagons stacked on top of each other the way graphite does, fullerenes can't be cut as easily as graphite.
    • However, fullerenes are very malleable, as the individual "hexagons" can rotate relative to one another.
    • Since each carbon atom is bonded to three others, fullerenes have a high melting and boiling point.
    • However, in fullerenes, each carbon forms a single bond with two other carbon atoms and a double bond with a third carbon atom. Therefore, fullerenes can't conduct electricity, as there are no delocalized electrons.

• Diagram of graphite's hexagonal pattern structure (works for fullerenes as well, but doesn't show the double bonds):

• Diagram of graphite's layers (doesn't work for fullerenes as fullerenes aren't arranged in layers this way:

• Diagram of diamond's structure:

Polarity

• Molecules have a property called polarity, which measures the net balance of electrons in the molecule.

• A molecule's polarity is due to its atoms' electronegativities. Usually, a molecule is made up of atoms with different electronegativities. Electronegativity is the tendency of an atom to attract electrons closer to its nucleus. When atoms share electrons in a covalent bond, the atom with more electronegativity will pull the shared electrons closer to itself, resulting in that atom having a slightly higher concentration of electrons. That atom will then be slightly more negatively charged than the other atom, which will be slightly more positively charged.

• This separation of charges is called a dipole. To measure how large the dipole is, find the difference between the electronegativities of the bonded atoms. The direction of the dipole is also shown by an arrow over the atoms that points in the direction of the more negatively charged atom, with a plus sign over the more positively charged atom. There is also a small delta sign underneath the atoms, with a positive or negative sign to denote which atom in the bond has a more positive or negative charge.

• Bonds can be distinguished by their dipoles:

  • Non-polar covalent bonds happen when the dipole between two bonded atoms is less than 0.5. There isn't much of a noticeable difference in the relative charges of the atoms.
  • Polar covalent bonds happen when the dipole between the two bonded atoms is between 0.5 and 1.8. The more electronegative atom pulls the shared electrons closer towards itself, but the electrons remain shared between the two atoms. The more electronegative atom gets a slightly more negative charge, while the other atom gets a slightly more positive charge.
  • Ionic bonds happen when the dipole between the two bonded atoms is greater than 1.8. The more electronegative atom just steals the shared electron$s$ from the less electronegative atoms. The more electronegative atom gets a negative charge, while the less electronegative atom gets a positive charge.

• A molecule's polarity is the combination of the dipoles of all the bonds between the atoms of that molecule. Dipoles that point in opposite directions will cancel each other out $completely if they have the same magnitude as well$. Usually, you'll only be asked to tell whether the molecule is polar or not $is the net polarity greater than 0.5 or not?$.

Electron & Molecular Geometry

• Electron geometry is the positioning of valence electrons in an atom. There are specific ways and positions in which they bond, which can be found through the VSEPR $Valence Shell Electron Pair Repulsion$ Theory. The valence electrons repel each other as much as possible, getting as far away from each other as possible, and in the process assume certain positions.

• The electron geometry is determined by the number of valence electrons there are - that determines how many atoms the central atom can bond with, which gives the electron geometry. Once the number of lone pairs is also known on the central atom, the molecular geometry can also be found.

Sigma & Pi Bonds

• Atoms can't directly form bonds to create molecules - their valence shells have multiple subshells $unless we're talking about Hydrogen$. These subshells have different energies, therefore their valence electrons have different energies.

• To get their valence electrons to have the same energies, the valence subshells hybridize to an intermediate energy spⁿ level. n is the number of p orbitals involved in the hybridization. Unless a p orbital is required to remain unhybridized so it can make a pi bond, it will hybridize. Sometimes, d orbitals will hybridize, but d orbitals will prefer to remain unhybridized unless they have to hybridize.

  • Hybridized orbitals all have the same energy, and thus repel each other equally, giving the Electron Geometry configurations.

• In a sigma bond, a hybridized orbital from one atom overlaps with a hybridized orbital from another atom to form a bond. If there are two electrons in a hybridized orbital, it doesn't have enough space for more electrons, and therefore can't bond. It is a lone pair.

• Since hybridized orbitals repel each other equally, and the nuclei also repel each other, each pair of atoms can only have one sigma bond, as the other hybridized orbitals aren't in the correct positions $assuming we're talking about the same two atoms. An atom that has a sigma bond with one atom can still have a sigma bond with another atom, it just can't have a second sigma bond with the same atom that it has already formed a bond with$.

• Any double or triple bonds are pi bonds, formed by the overlap of parallel unhybridized p orbitals that exist in a plane perpendicular to the hybridized orbitals.

• When the hybridized orbitals form the sigma bond, if there is a pi bond to also be formed, the p orbitals aren't close enough to be touching. The nucleus of each atom attracts the other atom's p orbitals, causing it to expand. The two expanding p orbitals meet and overlap in the middle, creating a pi bond.