CHEMICAL BONDING: 2a. Introduction to ionic bonding, ions, periodic table patterns & which elements are likely to combine to give an ionic compound

(1) IONIC BONDING - what are ions?

Ionic bonds are formed by one atom transferring electrons to another atom to form positive and negative ions which will attract each other due to being opposite in charge.

An ionic bond is most likely to be formed when a metal combines with a non–metal to form an ionic compound.

Elements consist of neutral atoms or molecules, the electrical neutrality is because the number of positive protons equals the number of surrounding negative electrons in their respective energy levels.

Charged particles called IONS are atoms, or groups of atoms, which have lost or gained one or more electrons to have a overall net electrical positive charge or negative charge.

In losing or gaining electrons to form an ion, part of an ionic bond, producing an ionic compound ...

the atoms try to attain a stable electron arrangement of a noble gas e.g. a full outer shell of electrons,

the number of electrons lost or gained gives the numerical charge of the ion,

an atom (usually non–metal) with a nearly full outer shell will try to gain electrons to form a negative ion,

the ion charge of –, 2–, 3– etc. results from the neutral atom gaining of 1, 2 or 3 electrons etc,

an atom (usually a metal) with a nearly empty outer shell will tend to lose electrons to form a positive ion,

the ion charge of +, 2+ or 3+ etc. results from the neutral atom losing 1, 2 or 3 electrons etc,

 and it is the mutual attraction of these positive and negative ions that will constitute the ionic bond.

Elements cannot form an ionic material i.e. an ionic bond can only be formed between two different atoms - usually a metallic element combining with a non-metallic element.

The atom losing electrons forms a positive ion (a cation) and is usually a metal.

The overall charge on the ion is positive due to excess positive nuclear charge (proton numbers do NOT change in chemical reactions) e.g.

Group 1 alkali metals lose their single outer electron to form single positive ions

e.g. Na ==> Na⁺ + e^–

Group 2 metals lose their two outer electrons to form doubly charged positive ions

e.g. Mg ==> Mg²⁺ + 2e^–

The atom gaining electrons forms a negative ion (an anion) and is usually a non–metallic element.

The overall charge on the ion is negative because of the gain, and therefore excess, of negative electrons e.g.

Group 7 halogen atoms gain one electron to form a singly charged negative ion

e.g. Cl + e^– ==> Cl^–

Group 6 non–metals gain two electrons to form a doubly charged negative ion

e.g. O + 2e^– ==> O^2–

(2) What is an ionic bond?

An IONIC BOND IS THE FORCE OF ATTRACTION BETWEEN ADJACENT IONS OF OPPOSITE CHARGE.

BUT, which electronic structures are the most stable to give ions that can form an ionic bond?

because this is what atoms will try to get to electronically!

  symbol (atomic number) electron arrangement

When atoms LOSE OR GAIN ELECTRONS, they try to attain the electron structure (electron configuration) of the electronically very stable atoms of the Group 0 Noble Gases

eg helium (2), neon (2.8) or argon (2.8.8) etc. quite simply because these are the most stable electron arrangements with a full outer shell of electrons (full highest energy level).

In advanced level chemistry you will encounter examples of electronic structures of ions that are NOT those of a Noble Gas.

Its a good idea to have some idea of where the elements are in the periodic table, and their electronic structure, before looking at the theoretical electronic model for ionic bonding in ionic compounds

work out the electron structure of simple ions

draw the electronic diagram of an ionic compound

draw dot and cross diagrams for ionic compounds formed by metals in Groups 1 and 2 with non–metals in Groups 6 and 7

work out the charge on the ions produced by metals in Groups 1 and 2 and by non–metals in Groups 6 and 7 and relate the charge to the group number of the element in the periodic table

deduce that a compound is ionic from a diagram of its structure in several different forms

describe the limitations of using dot and cross, ball and stick, two and three–dimensional diagrams to represent a giant ionic structure

work out the empirical formula of an ionic compound from a given model or diagram that shows the ions in the structure.

(3) Which element combinations from the periodic table produce ionic bonds and therefore an ionic compound?

The black zig–zag line 'roughly' divides the metals on the left from the non–metals on the right of the elements of the Periodic Table.

The electronic structures of the first 20 elements of the Periodic Table You need to know about these to understand the details of ionic chemical bonding

Pd metals Part of the modern Periodic Table Pd = period, Gp = group metals => non–metals Gp1 Gp2 Gp3 Gp4 Gp5 Gp6 Gp7 Gp0 1 1H  Note that H does not readily fit into any group 2He 2 3Li 4Be atomic number Chemical Symbol eg 4Be 5B 6C 7N 8O 9F 10Ne 3 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar 4 19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr 5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe 6 55Cs 56Ba Transition Metals 81Tl 82Pb 83Bi 84Po 85At 86Rn Gp 1 Alkali Metals,  Gp 2 Alkaline Earth Metals,  Gp 7 Halogens,  Gp 0 Noble Gases Ionic chemical bonding comments about the selected elements highlighted in white e.g. When the metals on the left, highlighted in white, combine with the non–metals, highlighted in white, on the right of the periodic table, an ionic bond is formed e.g. the formation of an ionic compound like sodium chloride NaCl or magnesium oxide MgO ion charge and group number Note: On this page to form stable ions with a noble gas electron arrangement occur by electron transfer ... For (a) to d) below M represents a group 1/2 metal and X represents a group 6/7 non-metal (a) Group 1 metals lose their 1 outer electron to form a singly charged positive ion: M ==> M+ + e– (b) Group 2 metals lose their 2 outer electrons to form a doubly charged positive ion: M ==> M2+ + 2e– (c) Group 6 non–metals gain 2 electrons to form a doubly charged negative ion: X + 2e– ==> X2– (d) Group 7 halogen non–metals gain 1 electron to form a singly charged negative ion: X + e– ==> X– In a correct ionic formula: total positive ion charge = total negative ion charge therefore we can predict the following formula where  M = a group 1/2 metal  and  X = a group 6/7 non–metal: (a) Group 1  +  (c) Group 6  ===>  M2X  or  (M+)2X2– (a) Group 1  + (d) Group 7  ===>  MX  or  M+X– (b) Group 2  +  (c) Group 6  ===>  MX  or  M2+X2– (b) Group 2  +  (d) Group 7  ===>  MX2  or  M2+(X–)2 Note that these formulae are empirical formula - the simplest whole number ratio of atoms in the formulae and is the ratio of the ions in their giant ionic lattice.

All the atoms of the metallic elements on the left–hand side of the Periodic Table eg (Groups 1/2) have 1/2 electrons in their outer shell (highest energy level) which are readily lost to form a positive ion of charge +1/+2 (cations) eg sodium, potassium, magnesium and calcium etc. The electronic structure of these stable positive ions are those of a Noble Gas with a full outer shell.

eg Na [2.8.1] ==> Na⁺ [2.8] like neon + e^–,

 or Ca [2.8.8.2] ==> Ca²⁺ [2.8.8] like argon + 2e^–

The atoms of the non–metallic elements on the right–hand side of the Periodic Table eg (Groups 6/7) have 6/7 electrons in their outer shell and try to gain 2/1 electrons to become electronically stable like a Noble Gas with a full outer shell of electrons eg oxygen and sulfur in Group 6 and the Group 7 Halogens – fluorine, chlorine, bromine and iodine. The electronic structure of these stable negative ions are those of a Noble Gas with a full outer shell.

eg O [2.6] + 2e^– ==> O^2– [2.8] like neon,

or Cl [2.8.7] + e^– ==> Cl^– [2.8.8] like argon

Brief summary of the Periodic Table including electronic structure and formula patterns

The examples below involve combining a metal from Groups 1 (Alkali Metals), 2 or 3, with a non–metal from Group 6 or Group 7 (The Halogens). The electron structures are shown in () or []. Only the outer electrons of the original atoms, and where they end up in the ions, are shown in the dot and cross (ox) diagrams

Ionic bonding is not directional like covalent bonding, in the sense that the force of attraction between the positive ions and the negative ions act in every direction around the ions which will alternate between being positive and negative to maximise attraction.

Note: In the examples of ionic bonding it is assumed YOU can work out the electron configuration (arrangement in shells or energy levels) given the atomic/proton number of the element from the Periodic Table.

See the main ionic bonding index for lots of dot and cross diagrams i.e. Lewis diagrams of ionic (electrovalent) bonding in ionic compounds

There are lots of dot and cross diagrams including simplified Lewis diagrams of ionic (electrovalent) bonding in ionic compounds

Lewis diagrams are quite minimalist.

For positive metal ions which give a noble gas structure when the electrons are 'lost' from the original metal atom, no outer shell electrons are shown!

For most simple non–metal negative ions, only the complete octet of outer shell electrons is shown for each atom.

Sodium chloride is a good example to start with.

Na ==> Na⁺ + e^–  (electron change from 2.8.1 to 2.8 like neon)

Cl + e^– ==> Cl^–  (electron change from 2.8.7 to 2.8.8 like argon)

and the simplest dot & cross diagram is shown below.

ONE atom combines withONE atom

to form

Note

(i) Ionic bonds are formed by electron transfer from a metal element to a non-metal element.

(i) The dot and cross electrons are identical, it just a way showing the origin of the electrons.

(ii) Only the outer shell of electrons is shown for the chloride ion, and no inner shells are shown for both ions.

There are several styles of diagrams - see ionic bonding in sodium chloride.

Note: Limitations of dot and cross electronic diagrams of ionic compounds

It is important to appreciate that dot and cross Lewis diagrams for ionic compounds do not show the structure of the compound in terms of ...

(i) arrangement of the ions in a crystal lattice of an ionic compound

(5) What is a crystal lattice - a giant lattice in terms of the ionic bonding in a compound?

IONIC BONDING – an ionic bond is formed by one atom transferring electrons to another atom to form oppositely charged particles called ions which attract each other – this electrostatic attraction is called an ionic bond and is most likely formed when a metal combines with a non-metal.

• An ion is an atom or group of atoms carrying an overall positive or negative electric charge

  • The electric charge is shown as a superscript +, –, 2+, 2– or 3+ etc.

  • e.g. Na⁺, Cl^–, [Cu(H₂O)₆]²⁺, SO₄^2– etc.

• If a particle, as in a neutral atom, has equal numbers of protons (+) and electrons (–) the overall particle charge is zero i.e. no overall electric charge.

• The proton/atomic number in an atom does not change BUT the number of associated electrons can!

• If negative electrons are lost the excess charge from the protons produces an overall positive ion.

• If negative electrons are gained there is an excess of negative charge, so a negative ion is formed.

• The charge on the ion is numerically related to the number of electrons transferred i.e. electrons lost or gained.

• For any atom or group of atoms, for every electron gained you get a one unit increase in negative charge on the ion, for every electron lost you get a one unit increase in the positive charge on the ion.

• The atom losing electrons forms a positive ion (cation) and is usually a metallic element.

• The atom gaining electrons forms a negative ion (anion) and is usually a non–metallic element.

• The ionic bond then consists of the attractive force between the positive and negative ions in the structure.

• The ionic bonding forces act in all directions around a particular ion, it is not directional, as in the case of covalent bonding.

•  

• The sodium (metal) atom transfers an electron to the chlorine (non–metal) atom in forming the ionic compound sodium chloride

• The bonds between the ions is very strong and they club together to form a giant ionic lattice with a very high melting point because it takes a lot of energy to overcome the attractive forces between the ions - the ionic bonds.

• When molten, or dissolved in water, ionic compounds will conduct electricity because the charged particles (ions) are free to move and carry the electric current.