4. Chemical bonding and structure

4. Chemical bonding and structure

4.1 Ionic bonding and structure

Nature of science:

• Use theories to explain natural phenomena - molten ionic compounds conduct electricity but solid ionic compounds do not. The solubility and melting points of ionic compounds can be used to explain observations.

Understandings:

• Positive ions (cations) form by metals losing valence electrons.

• Negative ions (anions) form by non-metals gaining electrons.

• The number of electrons lost or gained is determined by the electron configuration of the atom.

• The ionic bond is due to electrostatic attraction between oppositely charged ions.

• Under normal conditions, ionic compounds are usually solids with lattice structures.

Applications and skills:

• Deduction of the formula and name of an ionic compound from its component ions, including polyatomic ions.

• Explanation of the physical properties of ionic compounds (volatility, electrical conductivity and solubility) in terms of their structure.

4.2. Covalent bonding

Nature of science:

• Looking for trends and discrepancies - compounds containing non-metals have different properties than compounds that contain non-metals and metals.

• Use theories to explain natural phenomena - Lewis introduced a class of compounds which share electrons. Pauling used the idea of electronegativity to explain unequal sharing of electrons.

Understandings:

• A covalent bond is formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei.

• Single, double and triple covalent bonds involve one, two and three shared pairs of electrons respectively.

• Bond length decreases and bond strength increases as the number of shared electrons increases.

• Bond polarity results from the difference in electronegativities of the bonded atoms.

Applications and skills:

• Deduction of the polar nature of a covalent bond from electronegativity values.

4.3 Covalent structures

Nature of science:

• Scientists use models as representations of the real world - the development of the model of molecular shape (VSEPR) to explain observable properties.

Understandings:

• Lewis (electron dot) structures show all the valence electrons in a covalently bonded species.

• The “octet rule” refers to the tendency of atoms to gain a valence shell with a total of 8 electrons.

• Some atoms, like Be and B, might form stable compounds with incomplete octets of electrons.

• Resonance structures occur when there is more than one possible position for a double bond in a molecule.

• Shapes of species are determined by the repulsion of electron pairs according to VSEPR theory.

• Carbon and silicon form giant covalent/network covalent structures.

Applications and skills:

• Deduction of Lewis (electron dot) structure of molecules and ions showing all valence electrons for up to four electron pairs on each atom.

• The use of VSEPR theory to predict the electron domain geometry and the molecular geometry for species with two, three and four electron domains.

• Prediction of bond angles from molecular geometry and presence of non-bonding pairs of electrons.

• Prediction of molecular polarity from bond polarity and molecular geometry.

• Deduction of resonance structures, examples include but are not limited to C₆H₆, CO₃²⁻ and O₃.

• Explanation of the properties of giant covalent compounds in terms of their structures.

4.4 Intermolecular forces

Nature of science:

• Obtain evidence for scientific theories by making and testing predictions based on them - London (dispersion) forces and hydrogen bonding can be used to explain special interactions. For example, molecular covalent compounds can exist in the liquid and solid states. To explain this, there must be attractive forces between their particles which are significantly greater than those that could be attributed to gravity.

Understandings:

• Intermolecular forces include London (dispersion) forces, dipole-dipole forces and hydrogen bonding.

• The relative strengths of these interactions are London (dispersion) forces < dipole-dipole forces < hydrogen bonds.

Applications and skills:

• Deduction of the types of intermolecular force present in substances, based on their structure and chemical formula.

• Explanation of the physical properties of covalent compounds (volatility, electrical conductivity and solubility) in terms of their structure and intermolecular forces.

4.5 Metallic bonding

Nature of science:

• Use theories to explain natural phenomena - the properties of metals are different from covalent and ionic substances and this is due to the formation of non-directional bonds with a “sea” of delocalized electrons.

Understandings:

• A metallic bond is the electrostatic attraction between a lattice of positive ions and delocalized electrons.

• The strength of a metallic bond depends on the charge of the ions and the radius of the metal ion.

• Alloys usually contain more than one metal and have enhanced properties.

Applications and skills:

• Explanation of electrical conductivity and malleability in metals.

• Explanation of trends in melting points of metals.

• Explanation of the properties of alloys in terms of non-directional bonding.