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Intramolecular Forces

Intramolecular Forces

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Dative Covalent Bonds (Co-ordinate Bonding)

In a standard covalent bond there are a pair of electrons shared between two atoms where each atom donates one ...
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Metallic

Metallic bonding is the bonding of metal atom to another metal atom Unless a non-metal atom is present, a metal ...
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Covalent

Covalent forms between pairs of non-metal atoms Non-metal atoms need to receive electrons to fill the spaces in their outer ...
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Ionic

Ionic bonding is the bonding of ions that are held together through electrostatic attraction Ionic bonding occurs between metals and ...
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Dative Covalent Bonds (Co-ordinate Bonding)

Dative Covalent Bonds (Co-ordinate Bonding)

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  • In a standard covalent bond there are a pair of electrons shared between two atoms where each atom donates one electron
  • In a dative covalent bond there is a pair of electrons shared between two atoms however donated by one atom providing both electrons
  • In a coordinate/dative covalent bond:
    • The atom that accepts the electron pair is an atom that does not have a filled outer main level of electrons (electron deficient atom)
    • Atom that is donating the electrons has a pair of electrons that not being used in a bond (lone pair)
  • They are represented by an arrow which points towards the atom that is accepting the electron pair
  • This is only to show how the bond is made

e.g.

  • All co-ordinate bonds have exactly the same strength and length as ordinary covalent bonds between the same pair of atoms
Metallic

Metallic

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  • Metallic bonding is the bonding of metal atom to another metal atom
  • Unless a non-metal atom is present, a metal atom cannot donate electrons
  • The outer electrons of all metals are delocalised
  • Metals are a lattice of positively charged ions that are in a ‘sea’ of electrons

    Properties of Metallic Bonding

  • Good conductors of electricity (the delocalised electrons can move throughout the structure, therefore the electrons from the negative side of a terminal can join the electron sea at the one end (e.g. in a wire) whilst simultaneously a different electron can leave the wire at the positive terminal
  • Good conductor of heat
  • Metals have great strength as a result of the charge on the ions, a greater charge means the number of delocalised electrons is greater too. The stronger electrostatic attraction between the positive ions and the electrons. As well as the size of the ions, the smaller the ion, the closer the electrons are to the positive nuclei and thus the bond is stronger
  • Metals are malleable and ductile, as after a small distortion, each metal ion is exactly in the same environment as before, and therefore the shape is retained
  • Metals have a high melting point as a result of their large structure. There is a strong attraction between the metal ions and the delocalised sea of electrons.

    • Covalent

    Covalent

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    • Covalent forms between pairs of non-metal atoms
    • Non-metal atoms need to receive electrons to fill the spaces in their outer shell
    • The atoms share some of their outer electrons so that each atom has a stable noble gas arrangement

      Structure of Covalent Bonds

      Atoms with covalent bonds are held together through electrostatic attraction between the nuclei and the shared electrons

      e.g. In a hydrogen molecule (H2), the two protons are held together by the pair of electrons.

      Forming Covalent Molecules

    • a group of covalently bonded atoms is called a molecule
    • in the example of Chlorine (Cl), which exists as a gas and therefore is a molecule Cl2
    • Chlorine has 8 electrons (in the orbitals [Ne] 2s2 2p5), the 2 atoms share one pair of electrons and therefore produces a stable noble gas.
    • The molecule has no charge as there are no electrons have been transferred from one atom too the other

    • Double Covalent Bonds

    • In a double bond, 4 electrons are shared
    • Double covalent bonds are presented as: =
    • They are present in alkenes and molecules such as O2 (O=O)

    • Properties of Covalent Bonds

    • Low melting points (the intermolecular covalent bond although strong, the intramolecular forces are weak)
    • Poor conductors of electricity (as each molecule has no overall charge)

    • Co-ordinate (dative) Bonding

    • A standard covalent bond is the sharing of a pair of electrons. Usually, one electrons comes from one atom, and one from the other
    • In co-ordinate bonding, the electron pair is provided by only the one atom
    • The atom that accepts the electron pair is the atom which has an incomplete outer main level of electrons. (or electron deficient)
    • The atom that is donating the electrons has a pair of electrons that is not being used in the bonds, called a lone pair
    • Co-ordinate bonds are represented as an arrow, where the arrow points towards the atom which is accepting the electron pair.
    • A Co-ordinate bond is equal in strength and bond length to a regular covalent bond

    Ionic

    Ionic

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    • Ionic bonding is the bonding of ions that are held together through electrostatic attraction
    • Ionic bonding occurs between metals and non-metals
    • Electrons are transferred from the metal atom to the non-metal atom
    • Ions are formed when one or more electrons are transferred from one atom to another
    • The simplest ions are single atoms which have either lost of gained electrons which therefore have a full outer shell.
    • Electrostatic attraction holds positive and negative ions together. The result is a very strong bond.

    Forming ionic compounds

    • The ionic compounds are made up positively charged part and a negatively charged part
    • The overall charge of any compound is zero
    • All negative charges in the compound must balance the positive charge

    e.g. Na+  + Cl– -> NaCl (0 overall charge)

    Properties of Ionically Bonded Compounds

    • Always solids at room temperature
    • Giant structure
    • High melting point (in order to melt an ionic compound energy must be supplied to break up the lattice ions)
    • Conducts electricity well in dissolved in a solution or melted (as the ions that carry an electrical current are free to move in the liquid state but are not free in a solid state)
    • Brittle and shatter easily (the large lattice of alternating positive and negative ions, a physical collision of the material may cause these layers to move and have 2 positive ions touching which therefore repel)

    Shapes of Molecules

    Shapes of Molecules

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    Determining the Shapes of molecules: 

    The shape of the molecule can be found using the VSEPR model:

    e.g. CH3–

    1. Determine the central atom

    C

    1. Count the number of valence electrons

    C = +4

    1. Add one electron for each bonding atom

    +4 + 3 (hydrogens) = 7

    1. Add or subtract electrons for charge (see Top Tip)

    Negative ( – ) charge, therefore -1

    7 -1 = 6

    1. Divide the total of these by 2 to find the total

    number of electron pairs

    6 /2 = 3

    As CH3— has 3 bonding groups according to VESPR, and has no lone pairs, it is a Trigonal Planar.

    Permanent Dipole-Dipole

    Permanent Dipole-Dipole

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  • A molecule with a permanent dipole will have a weak electrostatic force of attraction between the δ+ and>δ– charges on the neighbouring molecules, otherwise known as a permanent dipole-dipole force
  • Permanent dipole-dipole is the 2nd strongest intermolecular force. It only acts between certain types of molecules
  • Molecules with a permanent dipole will experience dipole-dipole forces
  • It is found between molecules with a differing electronegativity, such as HCl. In HCl, the chlorine is more electronegative and thus will have a negative charge cloud around it (comparatively to the hydrogen atom) The result of this is that the strong negative chlorine will be attracted to the positive nucleus of the neighbouring Hydrogen in the HCl molecule.

    Dipole moments

    • Molecules that have polar bonds may also have a dipole moment
    • In molecules where there are more than one polar bond, the result of this is that the bonds may cancel each other out and thus there is no dipole moment. The effect also may add up and therefore each individual bond reinforces each other
    • It is greatly determined on the shape of the molecule

    • Van der Waals

    Van der Waals

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    • Van Der Waals is the weakest of all the intermolecular forces
    • All atoms and molecules have a positive and a negative charge despite being neutral overall
    • These charges result in the electrostatic attraction between all atoms and molecules, otherwise known as Van der Waals (VdW)
    • Larger molecules have larger electron clouds and therefore can create a larger dipole difference meaning there is a stronger Van der Waals
    • the shape of the molecule can also effect the Van der Waals, for example a long straight molecule can lie alongside another similar molecule thus the area of attraction is greater. This differs to a branched molecule as there is less surface area for attraction

    Forming Van der Waals

    • The random movement of electrons in their charge clouds means naturally for a moment of time one side of a molecule will have more electrons compared to the other
    • The molecule now has a temporary dipole
    • The shift in charge causes another neighbouring molecule to be attracted to it cause it to have a temporary dipole, however in the opposite direction (as electrons repel from each). The result is a positive and a negative end on a molecule, with the positive end being attracted to the temporary dipole negative end.
    • The chain continues with the 2nd dipole going on to effect other molecules in the same fashion
    • As the movement of electrons is random the dipoles are created and destroyed all of the time, despite this the molecules still remain attracted to one another
    Polarity

    Polarity

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    Polarity is the unequal sharing of electrons between 2 atoms which are covalently bonded together

    Non-Polar

    Non-polar molecules are ones that form between 2 non-metals which have the same electronegativity
    Covalent bonds in diatomic gases (e.g. O2) are non-polar as the atoms have an equal electronegativity and therefore the electrons are equally attracted to both nuclei
    Some elements such as carbon and hydrogen have similar electronegativities and so the bonds between them are non-polar

    Polar

    Polar molecules between 2 non-metal atoms that have a different electronegativity to one another and therefore have an unequal sharing of electron pairs
    If charge is distributed unevenly over a whole molecule, the molecule will have a permanent dipole
    Molecules with a permanent dipole are known as polar molecules
    In simple molecules, the one polar bond means the charge is distributed evenly across the whole molecule
    In complex molecules, several polar bonds may be present. The shape of the molecule will therefore be decided on whether or not there is an overall permanent dipole. If the polar bonds are arranged symmetrically so that the dipoles cancel each other out, then the molecule does not have a permanent dipole and is non-polar.
    If all of the polar bonds point in the same direction and do not cancel each other out. The resulting charge will be arranged unevenly across the entire molecule, resulting in a polar molecule
    The movement of electrons to the more electronegative atom can be displayed using an arrow: