Covalent Bonding & Ionic Bonding
Atoms are constantly striving for eight electrons in their outer shell; this is called the octet rule. In order achieve an outer shell of eight they must be willing to share or receive different amounts of electrons; valence is an atoms capacity to from bonds. Interaction between unpaired electrons and vacancies in energy shells are the cause of all atomic bonding. Ionic and covalent bonds are opposing ends of a spectrum of electrostatic interaction between atoms or ions. In reality there is a gradient from one to another.
Atoms are constantly striving for eight electrons in their outer shell; this is called the octet rule. In order achieve an outer shell of eight they must be willing to share or receive different amounts of electrons; valence is an atoms capacity to from bonds. Interaction between unpaired electrons and vacancies in energy shells are the cause of all atomic bonding. Ionic and covalent bonds are opposing ends of a spectrum of electrostatic interaction between atoms or ions. In reality there is a gradient from one to another.
Figure 1; triple covalent bond in nitrogen
Electrons interact in pairs– electrons negotiate their repelling charges via spin pair repulsion. Negatively charged covalence electrons may also reel electro-positive nuclei together. All bonds between any non-metallic substances will be covalent. A covalent bond is when neither atom has the strength of attraction to fully remove an electron from the other and orbiting electron pairs can be shared between two atoms. Nitrogen is not naturally found as an individual atom but is diatomic and occurs as an N2 molecule. Nitrogen has three valence electrons and five electrons in its outermost P-orbital. Nitrogen requires three in its outer shell to complete its octet. By sharing three electrons each. They now have six in their P-orbital - each nitrogen is allowed a temporary octet, including their inner 1s2 orbits. A dative covalent bond occurs when an electron pair is donated to an acceptor atom - dative covalent bonds are weaker.
Many simple covalent molecules have low boiling points e.g. nitrogen, ethanol, carbon dioxide and ammonia are all gasses at room temperature. N2 shares three pairs of electrons therefore N2 has three covalent bonds; each bond is named successively as s, p, and d. Three is the maximum amount of covalent bonds two atoms can share. It’s common for covalent substances to be water-insoluble; polymers and oils are covalent substances.
Figure 2; Ionic bonding in sodium chloride
Although covalent bonding is stronger – having a greater bond enthalpy than ionic bonding; ionic compounds are crystalline solids often with very high boiling points. An ionic bond forms when an electron is permanently removed and attached to another atom to form two oppositely charged ions. Oppositely charged ions are attracted to one another by electrostatic. Ionic bonds typically occur between metals and non-metals. Metals are very ready to pass on electrons; they are conductive and reduction agents. Non-metals are often oxidising agents and have a tight nuclear radius and therefore high electronegativity. When you react sodium and chlorine – sodium oxidizes to become a positively charged cation. Chlorine accepts an electron and is therefore reduced to from a negatively charged anion, forming NaCl. Sodium chloride is a brittle solid due to its ionic lattice. NaCl crystals, like many other ionic compounds (e.g. salts and acids) are soluble in water. Aqueous or molten ionic compounds are conductive.
Figure 3; Three-dimensional ionic lattice in sodium chloride
Intermolecular Forces
Circular visualisations of electron orbitals are a simplification. The actual path of an electron is defined by its orbital An orbital is in fact a cloud of the probable position of an electron, in effect the density of the cloud signifies where an electron is likely to be. An electron has the potential to deviate from its predicted position quite dramatically! Therefore in an atom electrons are not completely evenly distributed across the orbitals - where the majority of the electrons occupy a negative charge (d-) is induced; on the opposing side the proton creates a positive charge (d+); this is called a dipole. Dipoles can have a ‘domino effect’ by causing synchronisation of charge distribution among atoms in a close vicinity- this is called an instantaneous dipole is induced dipole. These weak intermolecular forces are also known as Van Der Waals forces.
Circular visualisations of electron orbitals are a simplification. The actual path of an electron is defined by its orbital An orbital is in fact a cloud of the probable position of an electron, in effect the density of the cloud signifies where an electron is likely to be. An electron has the potential to deviate from its predicted position quite dramatically! Therefore in an atom electrons are not completely evenly distributed across the orbitals - where the majority of the electrons occupy a negative charge (d-) is induced; on the opposing side the proton creates a positive charge (d+); this is called a dipole. Dipoles can have a ‘domino effect’ by causing synchronisation of charge distribution among atoms in a close vicinity- this is called an instantaneous dipole is induced dipole. These weak intermolecular forces are also known as Van Der Waals forces.
Dipoles can also be induced in molecules. Chlorine is diatomic; a chloride molecule is non-polar. Each chlorine atom has the same electronegativity therefore balanced forces of attraction allow for electrons to orbit impartially. There is an even distribution of electrostatic across their covalence orbital.
Figure 4; Polarity of a hydrogen-chloride molecule
Atoms have a force of attraction, this attracts electrons in a covalent bond, this called electronegativity. Electronegativity increases across a period and up through the groups – it would be fair to make the observation that it is related to atomic radius and levels of shielding, although evidently not completely due to irregularity in observed periodicity. In covalent compounds the constituent atoms are likely to have an imbalance in their electronegativity. This causes polarity in these molecules because the electrons will be biased to occupy one partnering atom over another. Hydrogen chloride HCL is a polar molecule – it has permanent dipole in which hydrogen is partially positively charged (d+) and chlorine is partially negative (d-). Similarly with Van Der Waals forces, permanent dipoles have influence on neighbouring particles; these are dipole-dipole interactions.
Hydrogen Bonds; Intermolecular Forces Continued
Hydrogen bonds are the strongest type of intermolecular forces and are essential for life on earth – they are present in water. Water is the medium in which all cell metabolism, interaction and biochemical synthesis is carried out. Hydrogen bonds are the present in a great deal of organic molecules.
A hydrogen bond is a relationship between a donor and an acceptor. The low electronegativity of hydrogen when bonded to a highly electronegative atom with a small physical radius - hydrogen’s proton is left almost completely unshielded, adopting a strong partially positive charge (d+). Hydrogen acts as a donor and is strongly attracted to unpaired electrons. Hydrides of nitrogen, oxygen and fluorine and that act as both an acceptor and a donor; intermolecular hydrogen bonds.
It’s important to understand that hydrogen bonds are at dipole-dipole interaction but are individualized because of their distinct properties. There are varying strengths of hydrogen bonds. Strong hydrogen bonds are more comparable to covalent or ionic bonds but are far weaker. Strong hydrogen dipole-dipole forces, strong hydrogen bonds occur in hydride anions in specialized conditions.
Moderate and weak hydrogen bonds are electrostatic attraction weaker than ionic. Moderate hydrogen bonds are the area of the most practical interest. The presence of moderate hydrogen bonds is reasonable for the unusual boiling point of water in comparison to other hydrides. Ice has a greater volume (therefore lower density) than water because hydrogen bonds organise oxidane particles into a lattice; hydrogen bonds are often represented by a dotted line. The strong polarisation of oxidane makes it a perfect solvent for many organic and inorganic reactions.
Figure 5; Two-dimensional visualisation of water molecules in a frozen lattice
Figure 7; Hydrogen bonding in two genetic base pairs (left guanine) (right cytosine)
By Caspar Zialor
Chemistry 1 by Brian Ratcliff, Helen Eccles, David Johnson, John Nicholson and John Raffan. 2003 | 08/11/2010 |
An Introduction To hydrogen Bonding by George A. Jeffrey 1997 |
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