Chemical Compounds

What are Chemical Compounds?

Chemical compounds are substances formed when two or more different elements chemically combine in fixed proportions.

• These compounds are held together by chemical bonds, which result from the interactions between the outer electrons of the atoms involved.
• Understanding the nature of these bonds and how atoms achieve stability through bonding is essential for predicting the properties and behaviour of compounds.

In most cases, atoms bond to attain a stable electronic configuration, often resembling that of noble gases.

• This process is driven by the tendency of atoms to achieve a full outer electron shell, which leads to the formation of either ionic or covalent bonds.
• The type of bond formed depends on the elements involved and their electron configurations, with bonding directly influencing the physical and chemical properties of compounds.

Valency

Valency refers to the combining power of an element, determined by the number of electrons an atom must lose, gain, or share to achieve a stable electronic structure.

Example: Carbon has a valency of 4 as it can form four covalent bonds to complete its outer shell (e.g., in $CH₄$).

Variable Valency of Transition Elements ($Cu$, $Fe$, $Cr$, and $Mn$)

Transition elements, especially those in the d-block of the periodic table, exhibit variable valency. This means they can form ions with different charges by losing varying numbers of electrons. The ability to adopt multiple oxidation states arises from the close energy levels of their outermost s and d orbitals, allowing different numbers of electrons to participate in bonding.

Here's a breakdown of the variable valencies of key transition metals: copper ($Cu$), iron ($Fe$), chromium ($Cr$), and manganese ($Mn$).

Copper ($Cu$)

• Valency: +1 and +2
• Explanation:
  • Copper can lose one electron to form the $Cu⁺$ ion, known as copper(I) or cuprous.
  • It can also lose two electrons to form the $Cu²⁺$ ion, known as copper(II) or cupric.
  • Example: $Cu₂O$ (copper(I) oxide) and $CuO$ (copper(II) oxide).

Iron ($Fe$)

• Valency: +2 and +3
• Explanation:
  • Iron can form the $Fe²⁺$ ion (iron(II), or ferrous) by losing two electrons.
  • It can also form the $Fe³⁺$ ion (iron(III), or ferric) by losing three electrons.
  • Example: $FeCl₂$ (iron(II) chloride) and $FeCl₃$ (iron(III) chloride).

Chromium ($Cr$)

• Valency: +3 and +6
• Explanation:
  • Chromium commonly forms $Cr³⁺$(chromium(III)), the most stable oxidation state in many compounds.
  • It can also form $Cr⁶⁺$(chromium(VI)) in compounds like chromates and dichromates, where it has a higher oxidation state.
  • Example: $Cr₂O₃$ (chromium(III) oxide) and $K₂Cr₂O₇$ (potassium dichromate).

Manganese ($Mn$)

• Valency: +2, +4, and +7
• Explanation:
  • Manganese can form $Mn²⁺$(manganese(II)), which is the most stable and common form.
  • It can also form $Mn⁴⁺$ in compounds like $MnO₂$ (manganese dioxide).
  • In its highest oxidation state, $Mn^{7+}$, it forms $MnO₄⁻$(permanganate ion), a powerful oxidising agent.
  • Example: $MnCl₂$ (manganese(II) chloride) and $KMnO₄$ (potassium permanganate).

Why Do Transition Metals Exhibit Variable Valency?

• Close Energy Levels: The s and d orbitals of transition metals are close in energy. Electrons from both orbitals can be lost or shared, allowing these metals to exhibit more than one valency.
• Stability of Oxidation States: The stability of different oxidation states depends on factors like the nature of the compound and the surrounding environment. Lower oxidation states (e.g., $Cu⁺$, $Fe^{2+}$) tend to dominate in ionic compounds, while higher states (e.g., $Fe³$⁺, $Mn⁷⁺$) are more common in covalent compounds or when the metal acts as an oxidising agent.

Stability of Noble Gas Electron Configurations

Noble gases (Group 18 elements) are characterized by their very stable electron configurations. Their outermost electron shells are fully occupied, making them highly unreactive under normal conditions. For example:

• Helium ($He$) has a full outer shell with 2 electrons.
• Neon ($Ne$), Argon ($Ar$), and other noble gases have a full outer shell with 8 electrons, adhering to the octet rule.

Uses of Helium and Argon Related to Their Chemical Unreactivity

Helium and argon, both noble gases, are chemically unreactive due to their stable electron configurations. This makes them highly suitable for various applications where reactivity could pose a problem. Their chemical inertness allows them to be used in situations requiring a non-reactive or protective environment.

Helium ($He$)

• Balloons and Airships: Helium is lighter than air and non-flammable, making it an ideal gas for filling balloons and airships. Unlike hydrogen, which is highly reactive and flammable, helium's chemical inertness ensures safety in these applications.
• Cryogenics: Helium is used as a coolant in cryogenics, particularly in superconducting magnets, such as those in MRI machines. Its low boiling point and non-reactivity make it suitable for cooling systems without causing chemical reactions or corrosion.
• Breathing Mixtures: In deep-sea diving, helium is used in breathing mixtures (such as heliox: helium and oxygen) to prevent nitrogen narcosis. Helium's unreactivity reduces the risk of adverse chemical reactions under high pressure.

Argon ($Ar$)

• Welding: Argon is widely used in arc welding and metal fabrication as a shielding gas. Its unreactive nature protects molten metals from reacting with oxygen and nitrogen in the air, which could weaken the weld or cause oxidation.
• Light Bulbs: Argon is used inside incandescent and fluorescent light bulbs to prevent the tungsten filament from oxidizing and burning out. Its chemical inertness ensures the filament can operate at high temperatures without reacting with the surrounding gas.
• Preservation of Historical Documents and Materials: Argon is sometimes used to create an inert atmosphere around sensitive historical documents, artworks, or metals to prevent them from reacting with oxygen or moisture in the air, which could cause degradation or corrosion over time.

Why Are Noble Gas Configurations Stable?

Full Outer Electron Shells

• Noble gases possess completely filled outer electron shells (s²p⁶ for all except helium), which provides them with maximum stability.
• This configuration has low energy, meaning noble gases do not readily form chemical bonds.

Minimal Reactivity

• Due to their full electron shells, noble gases have little tendency to gain, lose, or share electrons, which explains their inertness.
• For example, neon ($Ne$) and argon ($Ar$) do not typically form compounds.

Attainment of Stability in Other Elements

• Many elements react to achieve the same stable electron configuration as the nearest noble gas.
• For example, sodium (Na) loses one electron to form Na⁺, acquiring the electron configuration of neon (Ne), and chlorine (Cl) gains one electron to form Cl⁻, achieving the electron configuration of argon (Ar).

Importance in Bonding

The drive to attain the stability of noble gas configurations is the foundation of chemical bonding:

• Ionic bonding involves the transfer of electrons to achieve noble gas configurations (e.g., NaCl).
• Covalent bonding involves the sharing of electrons to complete outer shells (e.g., H₂O).

The Octet Rule

e.g. carbon and oxygen elements react to form carbon dioxide as

$$C + O2\space \rightarrow \space CO2$$

The noble gases in group 0 do not tend to react and form compounds. The reason for this is that they have stable configurations e.g. the main level electronic configurations of the first four noble gases are:

Noble gas element	Electronic arrangement
Helium (He)	2
Neon (Ne)	2,8
Argon (Ar)	2,8,8

Most other atoms react to try and achieve the stability of the noble gases. It is important to note that each of the noble gases except helium have 8 electrons in their outer level.

The octet rule states that when bonding occurs most atoms want to have 8 electrons in their outer level. This is a good working rule but there are exceptions e.g. lithium wants to lose one electron to have two electrons in its outer level like helium.

Using the Octet Rule to Predict the Formulas of Simple Compounds

The octet rule states that atoms tend to form bonds in such a way that they each attain eight electrons in their outermost shell, achieving a stable electronic configuration similar to that of the noble gases. This rule is useful for predicting the formulas of simple compounds, particularly for elements in the first 36 of the periodic table (excluding d-block elements).

Binary Compounds

Binary compounds are composed of two elements. By applying the octet rule, we can predict how many electrons each atom will gain, lose, or share to achieve a full outer shell.

Ionic Compounds:

• Formed between metals and non-metals.

• Metals lose electrons to form positive ions (cations), while non-metals gain electrons to form negative ions (anions). The ratio of ions in the compound reflects the need to balance their charges.

• Example: Sodium ($Na$), which loses 1 electron ($Na⁺$), combines with chlorine ($Cl$), which gains 1 electron ($Cl⁻$), to form sodium chloride ($NaCl$). Covalent Compounds:

• Formed between non-metals by sharing electrons to complete their outer shells.

• Example: In water ($H₂O$), oxygen shares two pairs of electrons with two hydrogen atoms, following the octet rule for oxygen and the duet rule for hydrogen.

Common Compounds Involving the Octet Rule

For elements in the first 36 of the periodic table, the octet rule helps predict formulas for several important types of compounds:

Hydroxides ($OH⁻$):

• Example: Sodium hydroxide ($NaOH$) is formed when sodium ($Na⁺$) combines with the hydroxide ion ($OH⁻$). Carbonates ($CO₃²⁻$):

• Example: Calcium carbonate ($CaCO₃$) forms when calcium ($Ca²⁺$) combines with the carbonate ion ($CO₃²⁻$), balancing the charges. Nitrates ($NO₃⁻$):

• Example: Potassium nitrate ($KNO₃$) forms when potassium ($K⁺$) bonds with the nitrate ion ($NO₃⁻$). Hydrogencarbonates ($HCO₃⁻$):

• Example: Sodium hydrogencarbonate ($NaHCO₃$), commonly known as baking soda, is formed by sodium ($Na⁺)$ and the hydrogencarbonate ion ($HCO₃⁻$). Sulfites ($SO₃²⁻$) and Sulfates ($SO₄²⁻$):

• Example: Magnesium sulfate ($MgSO₄$) forms when magnesium ($Mg²⁺$) combines with the sulfate ion ($SO₄²⁻$).

Predicting Formulas Using the Octet Rule

To predict the formula of a compound:

1. Determine the valency (or charge) of each element based on how many electrons are needed to achieve a full outer shell.
2. Balance the charges to ensure that the overall charge of the compound is neutral.

• For ionic compounds, the total positive and negative charges must cancel out.
• For covalent compounds, ensure the atoms share enough electrons to satisfy the octet rule.

Example Predictions:

• Calcium Chloride ($CaCl₂$): Calcium ($Ca$) has a valency of +2, needing two chloride ($Cl$) ions, each with a -1 charge, to form the neutral compound $CaCl₂$.
• Carbon Dioxide ($CO₂$): Carbon ($C$) needs four additional electrons, and each oxygen ($O$) atom needs two, so carbon shares two pairs of electrons with two oxygen atoms to form $CO₂$.

Exceptions to the Octet Rule:

While the octet rule is helpful for predicting formulas, there are exceptions:

• Hydrogen follows the duet rule, needing only two electrons.
• Elements in periods beyond the second, like phosphorus and sulfur, can have expanded valence shells, accommodating more than eight electrons.