Structure and Bonding
Molecules
What is a molecule?
Molecules are made of fixed numbers of atoms, joined together by covalent bonds and these molecules can vary in size from being single atoms (like in noble gases) or being very complex (like DNA, or polymers).
How would you describe a simple molecular structure?
Simple molecular substances are of molecules, whereby the atoms are joined together by strong covalent bonds.
However, between each covalent bond, the molecules are held together by weak intermolecular forces (London or Van der Waals forces).
These forces determine the boiling/melting point of the substance, and these type of structures typically have low boiling/melting points.
Simple molecules don’t conduct electricity as they do not have a sea of delocalized electrons, nor an overall electric charge.
Other examples of simple molecular substances are hydrogen (H2), ammonia (NH3), and water (H20). These substances have very strong covalent bonds between the atoms, but weak intermolecular forces holding the molecule together. When the substances are heated to melt/boil, these weak intermolecular forces break, not the strong covalent bonds.
A CO2 molecule contains one carbon atom bonded with 2 oxygen atoms, and is a simple molecular structure. The few atoms are held together by strong covalent bonds. Between one covalent bond and another, there’s Van der Waals forces (IMFs) that break down easy as these forces in these structures are weak (the weaker, the lower the boiling/melting point.)
Properties of Simple Molecular Structures
Why are simple molecular structures generally gases and liquids at room temperature?
Simple molecular structures are generally gases and/or liquids at room temperature as they have weak intermolecular forces of attraction.
Molecules have covalent bonds that hold the molecule together very strongly, but the forces are irrelevant to the physical properties of the substance (state of substance, boiling point, melting point, etc).
The physical properties of a substance is determined by the intermolecular forces (forces attracting one molecule to another molecule, like Van der Waals attraction or hydrogen bonds).
Simple molecular structures have comparatively weak intermolecular forces of attraction, and boiling/melting a molecular substance will not require breaking the covalent bonds the molecules have, as the melting/boiling point size depends on the strength of the intermolecular forces.
For example, if hydrogen bonds are present in the molecule, then the melting/boiling points of that substance will increase. Essentially, the greater the size of the molecule, there’s more van der Waals attractions formed, and to break these attractions/forces will require greater energy.
Giant Covalent Structure
Giant covalent structures contain many atoms joined together by covalent bonds to form a giant lattice.
They have high melting and boiling points. For example, Graphite and diamond have different properties because they have different structures.
Graphite conducts heat and electricity well because it also has free electrons.
Giant covalent structures are structures that contain multiple atoms joined together by strong covalent bonds to form and uphold a giant covalent lattice structure.
These structures have a significantly higher melting/boiling point compared to simple molecular structures.
Examples of Giant Covalent Structures
An example of a giant covalent structure is graphite and diamond.
These two giant covalent structures are allotropes, which allows their physical properties to differ from each other. Graphite can conduct heat and electricity well as it has free electrons within its structure. Diamond has a higher melting/boiling points than graphite.
Properties of Giant Covalent Structures
Substances with giant covalent structures have very high melting points, because a lot of strong covalent bonds must be broken, as these very strong bonds require great energy to overcome. For example graphite has a melting point of more than 3,600ºC.
Giant Covalent Structures (macromolecules) show variable conductivity. For example, diamond cannot conduct electricity, whereas Graphite can as it contains free electrons. Silicon is semi-conductive, where it’s half conductive and half not.
Giant Ionic Structures
A giant ionic structure is a complex and large metallic crystal structure comprised of positive (cations) and negative (anion) ions. These ions are arranged in a manner where the structure remains held together due to the opposite charges of the ions and the electrostatic force of attraction between them.
Diagram of Giant Ionic lattice structure: NaCl
Properties of Giant Ionic Structures
Giant ionic structures are able to conduct electricity when liquid or aqueous as the ions are able to move freely.
In a solid lattice, the ions are rigid and fixed in place, where they can vibrate but cannot move out of their position.
For example, if electrodes were put in a solid ionic compound, the ions would not be able to move towards them. If the ions are molten/aqueous, the ions become free to move around within the liquid, and therefore move towards the designated electrodes (if undergoing electrolysis).
The ions are attracted and move towards the oppositely charged electrode, carrying their charge through the liquid which allows a flow of electric current.
Metallic Bonding
Metallic bonding is where valence electrons (from the s and p orbitals) of the interacting metal atoms ‘delocalize’. The delocalization of electrons forms a “sea” of electrons that surround the positively charged nuclei of that metal ion.
Metallic Bonding in Magnesium atom
Properties of metals:
Metals are malleable:
Malleable means in can be hammered into sheets/shape, and metals have this property as their atoms have the ability of sliding over one another into new positions without breaking the metallic bond.
This is due to the metal atoms being of the same shape and size, meaning the arrangement of the atoms are not distorted, aiding in sliding layers over each other easily.
Metals can conduct electricity:
Metals are able to conduct as they have a “sea” of delocalised electrons, which then carry the charge of current through them.
The delocalised electrons (valence shell electrons from s and p orbital of metal atom) can freely move within the metal structure, and when an electrical current is applied, these electrons can carry the current.
Covalent Bonding
*Covalent bonds are formed between two nonmetal atoms, with a bond containing a shared pair of electrons between 2 positive nuclei.
Covalent bonds are formed when a non-metallic atom shares a pair of electrons with another non-metallic atom.
The non-metal covalent elements that form covalent bonds are Nitrogen, Hydrogen, Carbon, Phosphorous, Oxygen, Sulfur and Selenium.
Additionally, all the halogen elements (Fluorine, Chlorine, Bromine, Iodine and Astatine) and stable noble gases (Helium, Neon, Argon, Krypton, Xenon and Radon) are non-metal covalent elements.
These elements form bonds with one another by sharing electrons to form compounds.
Why is the covalent bond in an oxygen molecule a double bond?
A double covalent bond is formed where two pairs of electrons are shared between the atoms, rather than just one pair.
Two oxygen atoms are joined together with two pairs of electrons shared between the atoms rather than just one pair, thus being a double bond.
Diagram (dot and cross) of carbon dioxide molecule
Ionic bonding:
Ionic bonds are formed between a metallic and non-metallic atom.
By forming ionic chemical bonds, the atoms become more stable as they obtain full outer shells from the transfer of electron from one atom to another. In an ionic bond, the ions have the electronic structure of a noble gas (full outer shells).
The oppositely charged ions are arranged in giant ionic lattice structures due to the electrostatic attraction between oppositely charged ions, and these ions in a lattice are held together with very strong ionic bonds and a high temperature is required to melt the crystal.
Atoms losing electrons become positively charged ions, whereby the fixed number of protons in the nucleus is greater than the remaining negatively charged electrons in the shells.
Atoms that gain electrons become negatively charged ions, where the fixed number of protons in the nucleus is less than the number of negatively charged electrons in the shells.
When atoms are put in positions to lose or gain electrons, these atoms become ions and they have a net charge.
E.g: Charge of an Aluminium ion
An aluminium ion has a charge of +3.
Initially, the aluminium atom had a charge of +13, + (-13 from electrons =0). The metal atom charge became neutral due to the equal numbers of protons and electrons.
When an aluminium atom becomes an ion, it loses 3 electrons (being transferred to other atoms, making those atoms an ion as well), leaving behind 10 electrons.
This results in the charge of an aluminium ion being +3.
Charge of a Chloride ion:
A chloride ion has a charge of +1.
Initially, the charge of a chloride atom was neutral due to the equal numbers of protons and electrons cancelling the opposite charges out.
When a chloride atom becomes an ion, it gains 1 electron to complete its valence shell.
This results in the charge of a chloride ion being +1.
Dot and Cross diagram of Sodium Sulfide:
.
.
Electronegativity:
Unequal sharing of electrons result in the formation of polar bonds. These polar bonds occur when there’s a difference in the electronegativities of the bonded atoms, as the more electronegative atom exerts a greater pulling force on the shared pair of electrons.
An unsymmetrical bond, like a bent/v-shape bond (e.g H2O), is said to be polar. ‘Dipole’ is the term used to indicate that the bond type has two separate opposite electric charges.
Diagram of water molecule:
electronegativity is how strong an atom will attract to a bonding pair of electrons.
Delta negative (δ-) represents the more electronegative ion, and delta positive (δ+) is used to represent a partial charge (less than unit charge associated with ions)
Elements in a covalent molecule are situated as part of the non-metals (right hand side of periodic table) as they have electrons to share (not to lose).
The electronegativity value between these elements can show the relative polarity of the bonds, and the type of bond it forms.
Polarity:
The difference of electronegativity between two atoms can determine the relative polarity of bonds (compared with O2, which is two atoms of the same element).
Difference of electronegativity values (using Pauling scale of electronegativity)
Nonpolar covalent = electronegativity difference = 0 to 0.4
Polar covalent = electronegativity difference = 0.5 to 1.9
Ionic = electronegativity difference = 2.0 to 4.0
Polar Bonds:
Polar covalent
Polar covalent: A chemical bond between two atoms with low electronegativity difference, in which the electrons required to form a bond is shared unequally between two atoms.
The electrons from the bonded pair are attracted more towards the atom which is more electronegative (δ-).
Non-polar covalent
Non-polar covalent: Non-polar covalent bonds are bonds between 2 non-metal atoms that have the same electronegativity values, and therefore, equal sharing of the bonding electron pair as the difference in electronegativities = 0.
E.g: Water
In a polar covalent bond, the electrons shared by the atoms are (on average) closer to the Oxygen nucleus than the Hydrogen nucleus.
The electron pair is displaced toward the more electronegative atom.
This Oxygen atom then obtains a partial-negative charge while the less electronegative atom, Hydrogen, has a partial-positive charge.
Diagram of polar covalent bond:
E.g: H2
In H2, each H atom has an electronegativity value of 2.1.
The difference between 2 H atoms in a H2 molecule is 0, therefore the covalent bond between them is considered nonpolar.
Polar molecule
Polar molecule: A polar molecule is a molecule where the bond dipoles present do not cancel each other out, in which the bond dipoles present do not cancel each other out and thus results in a molecular dipole.
Cancellation depends on the shape of the molecule and the orientation of the polar bonds.
Molecular dipole
A molecular dipole: Bond dipoles may/may not cancel out producing either molecules that are nonpolar (if they cancel out) or polar (if they do not cancel out)
E.g: CO2 is a linear molecule with 2 bond dipoles that are equal and oppositely directed therefore the bond polarities cancel and the molecule is nonpolar.
Diagram of carbon dioxide bond:
Ionic bonds:
Ionic bonds: Ionic bonding is the complete transfer of valence electron(s) between atoms, that generates two oppositely charged ions.
The metal loses electrons to become a positively charged cation and the non-metal accepts those electrons to become a negatively charged anion.
Diagram of NaCl ionic bonding:
click to edit
Dative Covalent bonds:
Dative covalent bond:
A covalent bond is formed by two atoms sharing a pair of electrons, and held together because of the 2 positive nuclei attracting the electron pair.
A dative covalent or co-ordinate bond is a covalent bond where it involves a shared pair of electrons, but both electrons come from the same atom.
E.g: NH4
N has a lone pair of electrons in NH3, and this lone pair of electrons (from one atom) will bond with a H+ to form NH4.
Diagram of NH4 bonding:
Ions
Experiment:
How do we know ions exist?
A trivial method of demonstrating the presence of ions is in the process of electrolysis (in solution/molten state).
A solid ionic compound never conduct electricity., as their ions are too rigid to move freely, although they can vibrate.
Current is conducted through the molten/aqueous solution of an ionic compound, and subsequent discharge of ions in the designated electrodes is sufficient proof to suggest/show the existence of ions within a solution.
These ions are very small and the migration of ions in solution is never visible.
Polarisation as applied to ions
Polarisation as applied to ions is the distortion of the electron cloud (where electrons go around a nucleus) of one atom with another.
Bond symbols:
VSEPR Theory:
VSEPR is the valence shell electron pair repulsion theory.
This theory states that electron pairs located/found on the valence shell/ outer energy level of atoms repel each other and position themselves as far apart as the electrons possibly can.
The shape of the molecule is determined by the angles between the bonded atoms.
The total number of charge centers around the central atom determines the geometrical arrangement of the electrons.
Lone pairs of electrons have a higher concentration of charge than a bonding pair because the lone pairs are not shared between two atoms, so they cause more repulsion than bonding pairs.
Order of repulsion
Highest repulsion: Repulsion between two lone pairs.
Repulsion between lone pair and bonding pair
Lowest repulsion: Repulsion between two bonding pairs.
Shape of molecules (bond angles, bonding pairs, etc)
1) Molecules with two charge centres will position at 180° to each other (linear shape)
E.d: CO2
2) Molecules with 3 charge centres will position themselves at 120° to each other, (Planar triangular shape to electron distribution)
E.g: BH3
3) If there are lone pairs at the centre atom (rather than a bond), the angle will be less than 120°, to be 104.5° (Bent/V-shape)
E.g: H2O
4) Molecules with 4 atoms will position themselves at 109.5° to each other (tetrahedral shape to electron pairs)
E.g: CH4
5) If one or more of the charge centers are a non-bonding pair, the angle in which they position themselves will be less than 109.5°, to be 107° to each other. (Trigonal pyramidal)
E.g: NH3
6) Molecules with 5 charge centres will position them 120° to each other. (Trigonal bipyramidal)
E.g: PCl5
7) Molecules with six atoms are symmetrically arranged around the central atom, with 4 of the attachments positioned in a square plane with 90° bond angles.
The remaining two attachments are positioned perpendicular to the square plane at opposite ends of the central atom.
E.g: SF6
Summary of bonding:
The molecule polarity is dependent on the polar bonds it contains and its orientation (the shape of the molecule).
If the bonds are equal polarity and arranged symmetrically, their charge separations will oppose each other and cancel out.
The molecules carbon dioxide, bromine fluoride and methane are all non-polar because the dipoles cancel out.
Diagram:
If the bonds are not symmetrically arranged, then the polarities will not cancel out.
Any molecule with lone pairs will be polar unless the lone pairs cancel out with another lone pair.
E.g: hydrogen chloride, chloromethane (CH3Cl) and ammonia (NH3)
Intermolecular Forces:
Van der Waals forces: Weak intermolecular forces of attraction that occur when temporary dipoles (formed due to shifting electrons) cause induced dipoles in neighbouring atoms
Dipole-dipole forces: intermolecular forces between opposite charges on molecules with permanent dipoles.
Hydrogen bonding: a hydrogen bond is formed when a H atom bonded to N, O, or F is attracted to a lone pair of electrons on an atom of the neighbouring molecule.
Melting and Boiling points (with influence of intermolecular forces)
The stronger the intermolecular force, the higher the boiling/melting point of the substance, as these strong intermolecular bonds require very high energy to overcome them.
Van der Waals forces are the weakest form of intermolecular forces (low boiling point)
Dipole-dipole is stronger than Van der Waals forces.
Hydrogen bonding is the strongest form of intermolecular forces. (far greater boiling point)
Order of intermolecular forces from weakest to strongest:
London dispersion forces < dipole-dipole < Hydrogen bonding.
click to edit
8) There are 2 lone pairs of electrons on opposite sides of the central atom. The other four atoms connect to the central atom. The bond angles are 90° to each other. (square planar)
E.g: XeF4