The BE&HM Series ~ Part IV: Ionic Compounds, Elements & Oxidation State Defined

In the last post, I discussed the Group A, aka the Main Group, elements.  The chemical behavior of atoms  -- how they react/combine with other atoms in "the environment" -- is determined by the way electrons are configured orbiting about the nucleus.  All Main Group elements (with the biochemically relevant exception of hydrogen) share a valence shell (outermost shell) with 8 "vacancies" for electrons.  Having a filled valence shell with 8 electrons is the preferred energy state for these atoms -- this is called the Octet Rule -- and these atoms react/combine in order to satisfy this rule.  The atoms in each group (or column) contain the number of electrons in the group number, thus Na is in Group IA has 1e, C in Group IV has 4 e's, and Br in Group VII has 7 e's.

The Group IA and IIA elements lose electrons to form cations to meet the Octet Rule when the next lower filled shell now becomes the valence shell.  The Group VIA and VIIA elements have mostly filled shells and will gain electrons to completely fill them.

In this installment I'll discuss how different atoms combine to form unique substances called compounds.   One type of compound, ionic, will be discussed in detail, and the formation of an ionic compound shall be used to introduce the concept of oxidation state.

Elements & Compounds

Element:  A substance consisting of only one atom.  Several elements form molecules* composed of two atoms as their most stable state in nature.  These are  H, O, N and all Group VIIA elements.  

Compound:  Two or more different atoms "bonded" in fixed proportions and the same geometric arrangement. In this series, I'll discuss two types of compounds:

• Ionic compounds
• Covalent compounds

These two types of compounds are distinguished by the nature of the bonds holding the atoms together.  Covalent compounds are by far the most prevalent in biochemistry and these compounds exist as distinct particles we call molecules.  [*The elements that form molecules are not considered compounds because the atoms are the same.  These all form molecules of two atoms, called diatomic molecules.]   I'll devote several posts to the discussion of covalent compounds.  This post will focus on compounds formed when ions combine.

Review on Ions

• Atoms that easily form ions have either mostly empty or mostly full outer valence shells
• Atoms that lose one or more electrons are positively charged ions called cations
• Atoms that gain one or more electrons are negatively charged ions called anions
• Group IA lose 1e to become +1 cations; Group IIA lose 2e's to become +2 cations
• Group VIIA gain 1e to become -1 anions; Group VIA gain 2e's to become -2 anions
• All of these ions are called monatomic ions as they are formed from a single parent atom

A few new things:

• All metals -- left hand side (all the way across to that "staircase") of the periodic table including Group A & B elements -- have a tendency to form cations.  They may not form them or be very resistant to forming them, but they do not form anions.
• Non-metals -- upper right side of the periodic table excluding Group VIIIA -- form anions, and this is mostly from Group VIIA in biological contexts.
• Non-metals also form covalent compounds (next installment) and the biologically relevant N, P, O, and S atoms tend to bond in this fashion in biological molecules.
• Non-metals combine to form compounds that can be charged, called polyatomic ions.  You may recognize some of these such as hydroxide (OH^-), sulfate (SO4^-2) and phosphate (PO4^-3).   I'll discuss these ions subsequent to the discussion of covalent bonding.
  
  

  Ionic Bonds

Ionic bonding involves the transfer of electrons.  Electrons do not exist in stable form floating out there in a vacuum -- they like to be associated with atoms, or more specifically with the protons in the nuclei of atoms where their negative charges are balanced out by the positively charged protons.  Electrical neutrality is a preferred state of matter.  Let's begin by looking at the classic ionic compound everyone is familiar with, sodium chloride, NaCl, aka salt! 

Sodium in Group IA has a single valence electron that is highly vulnerable to being knocked off.  If all I have is pure sodium, those electrons have a high propensity to want to fly off, but nowhere to go.  This is why pure sodium is so highly reactive.   Chlorine is Group VIIA meaning it has one vacancy to fill to satisfy the Octet Rule.  In the diagram below we see the parent atoms with their electron configurations.  The lone electron in the M shell of Na finds a happy home in the Cl atom, Na loses an electron becoming  Na⁺ and Cl gains an electron becoming Cl^- .  These oppositely charged ions are attracted to one another like the opposite poles of two bar magnets are attracted to one another -- the ions then stick to one another just like the bar magnets would.

So an ionic bond is the electrostatic force between oppositely charged ions

Ionic Compounds

A compound composed of one or more type of cation and one or more type of anion held together by electrostatic forces.

Some important things about ions and ionic compounds:

• The formation of an ionic bond involves the TRANSFER of electrons from the cation parent atom to the anion parent atom. 
• An ionic compound requires at least one cation and one anion
• Metals form cations and non-metals form anions, so an ionic compound must contain at least one of each type of atom, metal and non-metal.
• Once this transfer occurs, the cations and anions are the more stable entities -- the cations don't want their electrons back, and the anions aren't giving back the electrons that made them "complete". Further electron movement between established I,II,VI and VIIA ions is not energetically favored.
• The electrostatic forces, ionic bonds, are relatively strong and most ionic compounds are solids at room temperature and have very very high melting points.
• The driving force in forming ionic compounds is to form a neutral substance, therefore cations and anions will combine in proportions so that the substance has a net charge of zero.
• There is no distinct "ionic molecule" -- ionic compounds exist as crystals that are formed depending on relative ionic sizes and proportions into different conficurations.
• The "unit" to describe an ionic compound is called the empirical formula and represents the lowest whole number proportions of ions that result in a neutral substance.
• An ionic compound is made up of charged particles, but overall is not a charged substance

So, sodium chloride is NaCl, a neutral compound with the crystal structure shown in the diagram above right.  No subscripts imply a 1:1 ratio which is to be predicted because Na loses 1 electron and Cl gains 1 electron, and +1 + (-1) = 0.   All IA and VIIA atoms (H excluded) can form ionic compounds at 1:1 ratios.   Now magnesium, Mg, is IIA so it forms +2 ions.  If it forms a compound with bromine, Br, it will form MgBr2 because each VIIA bromine can take one electron:  +2 + 2(-1) = 0.  But Mg will form 1:1 ionic compounds with VIA, for example magnesium oxide MgO. 

Oxidation & Reduction, Redox Reactions & Oxidation State

As mentioned previously, the formation of an ion from its parent atom involves losing or gaining electrons.   

• Oxidation is defined as the loss of electrons
• Reduction (the opposite of oxidation) involves gaining electrons

My students taught me this memory tool:  LEO goes GER  I'm not a big fan, but LEO is a lion that roars, GER, and it seems to help many of them.  LEO = Lose Electrons Oxidation, GER = Gain Electrons Reduction.   We can neither create nor destroy electrons, so:

For oxidation of one atom to occur, reduction must occur in another -- while we often focus on one or the other phenomenon, this must always hold.  

The oxidation state of an atom is a number that indicates the number of electrons associated with the atoms in excess or deficit of the atomic number.  It is sometimes called the oxidation number.  Recall the atomic number is the number of positively charged protons in an atom and also the number of negatively charged electrons in the neutral atom.   Some rules for determining oxidation state in atoms and ions:

• All pure elements have an oxidation state of zero, 0.  This includes pure elements that form diatomic molecules.  
• All monatomic ions have oxidation states equal to the charge on the ion.  For example, Na⁺ has an oxidation state of +1 and Cl^- has an oxidation state of -1.
• For ions, a positive oxidation state signifies a relinquishing of ownership of the electrons, a negative oxidation state signifies ownership of extra electrons.
• All group IA, IIA, VIA and VIIA elements will have the oxidation states of their ions when in ionic compounds.
• The Transition or Group B metals generally have more varied oxidation states depending on the compound in which they are a part.

Let's look at the reaction for the formation of sodium chloride from sodium and chlorine.  The YouTube video below is one of these "don't try this at home" things, but it demonstrates both the reactivity of sodium and its reaction with chlorine gas.  I'm not sure what he mixes in that cylinder to create the chlorine gas (a common one is calcium hypochlorite and hydrochloric acid) but he ignites the sodium to initiate the reaction when the sodium is brought into contact with the chlorine.  Unfortunately he doesn't clearly show the product, which is salt, NaCl.

The chemical reaction is shown at right.  This balanced stoichiometrically so that we see that we need 2 moles of sodium metal to react with each mole of the diatomic chlorine gas to form 2 moles of sodium chloride.  This reaction is a redox couple that can be broken down into half reactions as shown.  The electrons produced in the oxidation half reaction are consumed in the reduction half reaction.  Therefore, the reaction is also balanced in terms of the number of electrons.  Looking back to the original equation, on the left, the pure elements Na and Cl have oxidation states of zero, 0.  On the right the product is an ionic compound and we assign the oxidation states as the charges on the ions.  Therefore:

Oxidation = loss of electrons →  increase in oxidation state or oxidation number

Reduction = gain in electrons → decrease in oxidation state or oxidation number

A note about the Group B, Transition Metals that form cations using iron, Fe as an example.  Iron, like most transition metals, can form multiple stable ions.  The two most common are the ferrous ion, Fe⁺², and the ferric ion, Fe⁺³.   Some possible oxidation reactions for iron are:  

Fe → Fe⁺² + 2e   ,    Fe → Fe⁺³ + 3e     ,     Fe⁺² → Fe⁺³ + e    

Some possible reduction reactions involving iron are the reverse:

Fe⁺² + 2e → Fe    ,    Fe⁺³ + 3e → Fe    ,     Fe⁺³ + e → Fe⁺² 

Therefore, Fe⁺² can be either an oxidized form of iron if formed from Fe, or a reduced form of iron if formed from Fe⁺³ .   Some final notes on the generalizations I've made in this post.  

• While metals generally form cations, some are capable of having negative oxidation states
• Hydrogen has been excepted from all generalizations of Group IA as it is classed as a non-metal and is a small atom desiring a filled shell with 2e's.  Hydrogen will be discussed separately
• I have only discussed the behavior of the non-metals and oxidation states when part of ionic compounds.  I'll discuss them as part of covalent compounds if and when it is relevant.
• If you're interested, a listing of possible oxidation states for the elements can be found here.  The generalizations I've made explain apparent conflicts ... e.g. Cl has an oxidation state of -1 in ionic compounds.  I'll address the positive oxidation states if they are ever relevant to a biological molecule.

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Next up, Covalent Bonding