Module 2. Electrons, bonding and structure 6-9

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Ionic bonding

Ionic bonds
An ionic bond is the electrostatic attraction between oppositely charged ions. Ionic bonds are present in compounds consisting of a metal and a non-metal.
Imagine an ionic bond being formed between two atoms.

  • Electrons are transferred from the metal atom to the non-metal atom.
  • Oppositely charged ions are formed which are bonded together by electrostatic attraction.
  • The metal ion is positive (loses an electron)
  • The non-metal ion is negative (gains an electron) 

Giant ionic lattice

3-dimensional structure of oppositely charged ions, held together by strong ionic bonds

  • Each ion is surrounded by oppositely charged ions
  • These ions attract each other from all directions, forming a giant ionic lattice. 
  • Sodium Chloride, NaCl, is also an ionic compound and forms a giant ionic lattice. Each Na+ ion surrounds 6 Cl- ions, and Cl+ ion surrounds 6 Na+ ions.
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Ions and the periodic table

Predicting ionic charges
You can predict the charge on an element's ion from its position in the Periodic Table. You can find the number of electrons in the outer shell from the element's position in the Periodic Table. It is then a simple step to calculate how many electrons need to be lost or gained in order to reach a noble gas electron configuration. From this, you can predict the likely charge on the resulting ion.
For example:

  • lithium, in Group 1, has one electron on its outer shell

To form the electron configuration of the nearest noble gas, helium:

  • a lithium atom must lose one outer electron
  • a lithium ion, Li+, therefore has a charge of 1+ 

Elements in the same group of the Periodic Table have the same number of outer-shell electrons and react in similar ways.

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Ions and the periodic table

Atoms of metals in group 1 - 3

  • lose electron
  • form positive ions with the electron coniguration of the previous noble gas in the periodic table

Atoms of non-metals in groups 5 - 7:

  • gain electrons
  • form negative ions with the electron configuration of the next noble gas

Atoms of Be, B, C and Si:

  • do not normally form ions
  • too much energy is needed to transfer the outer-shell electrons to form ions

Some elements can form more than one ion, each with a different charge. The oxidation number of the element is written as a Roman numeral.

  • iron (II) for Fe2+ and irion (III) for Fe3+
  • copper (I) for Cu+ and copper (II) for Cu2+
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Ions and the periodic table

Molecular ions
Groups of convalent-bonded atoms can also lose or gain electrons to form ions. These are called molecular ions.

  • Molecular ions with 1+ charge: Ammonium NH4+ 
  • Molecular ions with 1- charge: Hydroxide OH-,Nitrate NO3-, Nitrite NO2-, Hydrogencarbonate HCO3-
  • Molecular ions with 2- charge: Carbonate CO3 2-, Sulfate SO4 2-, Sulfite SO3 2-
  • Molecular ions with 3- charge:Phosphate PO4 3- 
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Covalent bonding

A covalent bond is a bond formed by a shared pair of electrons.
A lone pair is an outer-shell pair of electrons that is not involved in chemical bonding.

Covalent bonds
Covalent bonds occurs in compounds consisting of non-metals. In a covalent bond, an electron pair occupies the space between the two atoms' nucleus.

  • The negatively charged electrons are attracted to the positive charges of both nuclei.
  • This attraction overcomes the repulsion between the two positively charged nuclei
    Bonding in hydrogen: two hydrogen atoms each share one electron ( 

We can imagine a covalent bond formed between two atoms of non-metals.

  • Two electrons are shared.
  • This is in contrast to the transfer of electrons that result in ionic bonds .
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Single covalent bonds

Hydrogen, H2, is formed from two hydrogen atoms.

  • Each hydrogen atom has one electron in its outer shell
  • Each hydrogen atom contributes one electron the the covalent bond

Bonding in hydrogen: two hydrogen atoms each share one electron (

A covalent bond is directional, acting solely between the two atoms involved in the bonds. Remember that an ionic bond attracts in all directions.
Remeber the number of covalent bonds formed by the atoms of these 4 elements:

  • Carbon (C) - 4 covalent bonds
  • Nitrogen (N) - 3 covalent bonds 
  • Oxygen (O) - 2 covalent bonds
  • Hydrogen (H) - 1 covalent bond
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Lone pairs

In a covalent bond, an electron pair is shared between two atoms, bonding them together. Sometimes though, the electron pair is not used for bonding. In this instance, the pair is known as a lone pair of electrons. In NH4 and H2O, an ammonia molecule has one lone pair and a  water molecule has two lone pairs.

A lone pair gives a concentrated region of negative charge around the atom. Lone pairs can influence the chemistry of a molecule in several ways. 

Multiple covalent bonds
Some non-metallic atoms can share more than one pair of electrons to form a multiple bond

  • Sharing of 2 pairs of electrons forms a double bond e.g. O2.
  • Sharing of 3 pairs of electrons forms a triple bond e.g. N2.
  • Carbon dioxide, CO2, has two double bonds 
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Further covalent bonding

Dative covalent bond
A dative covalent or coordinate bond is a shared pair of electrons which has been provided by one of the bonding atoms only.
In a dative covalent bond, one of the atoms supplies both the shared pair of electrons to the covalent bond.

  • A dative covalent bond can be written as A-B
  • The direction of the arrow shows the direction in which the electron pair has been donated 

The ammonium ion, NH4+
The ammonium ion, NH4+ has 3 covalent bonds and one dative covalent bond. One of the electron pairs around the nitrogen atom in an NH3 molecule is a lone pair. In the formation of an ammonium ion, this lone pair provides both the bonding electrons when bonding with the H+ ion, resulting NH4+ ion has a positive charge of 1+.

In a dative covalent bond, one atom provides both bonding electrons from a lone pair of electrons. However, once formed, this dative covalent bond is equivalent to all of the other covalent bonds. For example, the NH4+ ion, you cannot tell which bond was formed from the N lone pair. 

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Further covalent bonding

How many covalent bonds?
When covalent bonds form, unpaired electrons often pair up so that the bonded atoms obtain a noble gas electron configuration, obeying the Octet rule (the tendency to acquire a noble gas electron configuration). 
This is not always possible.

  • there may not be enough electrons to reach an octet
  • more than four electrons may pair up in bonding (expansion of the octet

Not enough electrons to reach an octet
Within period 2, the elements beryllium, Be, and boron, B, both form compounds with covalent bonds. However, Be and B do not have enough unpaired electrons to reach a noble gas electron configuration. But they can pair up any unpaired electrons. 

Expansion of the octet
For elements in group 5-7, something odd happens from period 3.
As we move done the Periodic Table, more of the outer-shell electrons are able to take part in bonding. In the resulting molecules, one of the bonding atoms may finish up with more than eigh electrons in its outer shell. This breaks the Octet Rule and is often called the expansion of the octet.  

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Further covalent bonding

Expansion of the octet (Example)
In the compound sulfur hexafluoride, SF6, sulfur has expanded its octet. 

  • Sulfur has 6 electrons in its outer shell
  • Six covalent bonds can be formed
  • Each sulfur's 6 electrons is paired
  • 12 electrons surroud S
  • Each of the six fluorine atoms have 8 electrons in its our shell, attaining the octet.

A better rule
A better rule than the Octet Rule would be:

  • unpaired electrons pair up;
  • the maximum number of electrons that can pair up is equivulant to the number of electrons in the outer shell 
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